LABORATOEY  METHODS 

OF 

INOEGANIC  CHEMISTEY 


BY 
HEINRICH   BILTZ  AND  WILHELM    BILTZ 

ri 

UNIVERSITY  KIEL  UNIVERSITY  GOTTINGEN 


U  N  AUTHORIZED  TRANSLATION   BY 
WILLIAM   T.  HALL  AND  ARTHUR   A.  BLANCHARD 

MASSACHUSETTS   INSTITUTE   OF  TECHNOLOGY 


NEW  YORK 

JOHN    WILEY   &   SONS 
LONDON:    CHAPMAN   &   HALL,  LIMITED 


^o^cxix. 


COPYRIGHT,  1909, 

BY 

WILLIAM  T.  HALL  AND  ARTHUR  A.  BLANCHARD 


Stanhope  ipreea 

F.   H.  GtLSON     COMPANY  G~29 

BOSTON.     U.S.A. 


AUTHORS'   PREFACE. 


THIS  book  outlines  a  course  of  laboratory  work  which  is  essen- 
tially synthetic  in  nature  and  is  designed  to  aid  in  acquiring  a 
more  adequate  knowledge  of  inorganic  chemistry  than  is  to  be 
obtained  by  practice  in  chemical  analysis  alone.  The  need  of 
supplementing  the  work  of  analytical  chemistry  in  such  a  way 
has  indeed  been  recognized  in  most  chemical  laboratories  of  the 
German  technical  schools  and  universities,  in  fact  such  instruction 
still  persists  as  a  part  of  the  classic  method  of  teaching  chemistry. 
To-day  a  training  according  to  these  lines  is  considered  de- 
sirable even  for  those  students  with  whom  chemistry  is  a  minor 
subject.  It  has  been  the  aim  of  the  authors  to  enlarge  the 
choice  of  suitable  experiments  by  publishing  the  procedures 
which  they  have  tested  and  found  satisfactory  in  their  own 
laboratories  in  the  course  of  a  number  of  years'  experience. 

The  experimental  part  of  the  book  is  given  in  relatively  com- 
plete detail,  because  in  our  opinion  this  does  not  materially 
lessen  the  self-reliance  of  the  student,  who  has,  as  a  rule,  but  a 
limited  time  at  his  disposal.  The  beginner  obtains  ample  oppor- 
tunity for  acquiring  manipulative  skill  and  for  exercising  inge- 
nuity, in  the  carrying  out  of  the  work  and  in  modifying  and 
improving  the  directions  to  correspond  to  the  facilities  at  hand; 
others  who  need  particular  preparations  as  starting  material  for 
further  investigation  can  appreciate  directions  in  which  technical 
difficulties  are  guarded  against. 

To  aid  in  the  study  of  the  theoretical  relations  brief  general 
discussions  are  interspersed  throughout  the  book  as  well  as  ref- 
erences to  the  original  literature  and  the  text-books  of  inorganic 
and  theoretical  chemistry  which  should  be  freely  consulted. 

As  regards  the  arrangement  of  the  material,  we  have  departed 
from  the  disposition  which  is  common  to-day  of  treating  the 
compounds  in  connection  with  the  groups  of  elements.  We  have 
chosen  to  base  the  classification  upon  the  different  types  of 

iii 


iy  AUTHORS'   PREFACE. 

combination,  and  thereby  have  returned  to  the  older  usage. 
Among  the  early  authors  Thenard  in  1813  writes:  "La  methode, 
que  fai  constamment  suivie,  consiste  a  proceder  du  simple  au  com- 
pose, du  connu  a  I'inconnu,  a  reunir  dans  un  meme  groupe  tons  les 
corps  analogues,  et  a  les  etudier  d'abord  d'une  maniere  generate  et 
ensuite  d'une  maniere  particuliere."  Gmelin  in  1817,  however, 
took  a  different  stand  and  in  his  "Handbuch"  arranged  the  com- 
pounds according  to  the  elements,  and  thus  introduced  the  sys- 
tem of  classification  which  has  been  followed  by  nearly  all  of  his 
successors.  Even  to-day  Gmelin's  reasons  for  departing  from 
the  older  method  of  treatment  probably  hold  equally  well  as 
regards  elementary  books  for  beginners;  the  older  method  scatters 
the  compounds  of  a  single  element  in  such  a  way  that  the  student 
fails  to  get  a  distinct,  coherent  picture.  This  book,  however,  is 
intended  primarily  for  those  who  have  passed  beyond  the  more 
elementary  stage  in  their  study  of  chemistry,  so  that  it  does  not 
seem  to  us  to  be  too  daring  to  make  the  experiment  as  to  how 
well  our  modern,  inorganic  chemistry  will  fit  into  the  older 
framework.  It  seems  as  if  thereby,  aside  from  the  old  advan- 
tage of  a  better  comprehension  of  analogous  methods  of  prep- 
aration and  analogous  properties,  there  results  a  particularly 
intimate  amalgamation  of  experimental  and  theoretical  chemistry, 
for  in  this  way  we  advance,  as  it  were,  from  a  "one  compo- 
nent system"  to  one  of  "several  components."  A  similar  return 
to  the  older  method  of  classification  is  to  be  found,  for  example, 
in  the  second  part  of  the  Modern  Chemistry  by  Sir  W.  Ramsey; 
in  the  arrangement  adopted  by  A.  Werner  and  P.  Pfeiffer  in  the 
Inorganic  Abstracts  in  R.  Meyer's  Jahrbuch  der  Chemie;  and  in 
A.  Werner's  Neuere  Anschauungen  auf  dem  Gebiete  der  anor- 
ganischen  Chemie.  This  system  was  also  outlined  by  one  of  us  in 
a  lecture  on  Complex  Compounds  delivered  in  Gottingen  in  1903- 
1904. 

The  method  of  arrangement  chosen  often  separates  prepara- 
tions which  are  closely  related  to  one  another,  as  for  example 
where  one  is  used  as  starting  material  for  the  formation  of  the 
other;  to  show  these  relations  the  "Dependent  Preparation" 
is  noted  at  the  end  of  many  of  the  procedures.  In  the  last 
chapter,  which  treats  of  compounds  of  the  rarer  elements,  we 
have  departed  from  our  chosen  system  of  classification  because 


AUTHORS'  PREFACE.  V 

in  this  case  the  question  of  raw  material  for  the  preparation  of 
the  individual  compounds  is  of  chief  importance,  and  here  the 
characterization  of  the  element  in  question  is  at  present  of  more 
importance  than  the  type  of  combination. 

The  instructor  will  at  once  notice  on  glancing  through  these 
pages  that  for  a  rapid  and  successful  carrying  out  of  the  pro- 
cesses certain  requirements  are  to  be  met  as  regards  apparatus 
and  laboratory  facilities;  these  should  at  all  times  be  available 
for  the  student's  use. 

The  experiments  necessary  for  working  out  and  testing  the 
methods  prescribed  were  carried  out  during  the  last  eight  years 
with  the  aid  of  students  at  the  University  of  Kiel,  and  to  some 
extent  at  the  Universities  of  Gottingen  and  Clausthal.  Published 
processes  have  frequently  been  changed  or  improved  according  to 
our  experience,  though  no  doubt  it  will  still  be  found  that  the 
directions  and  yields  can  be  further  improved. 


TRANSLATORS'  PREFACE. 


FOE  some  time  previous  to  the  appearance  of  the  German 
edition  of  this  book,  the  translators  had  been  convinced  of  the 
value  of  laboratory  work  in  Inorganic  Preparations  as  a  basis  for 
the  teaching  of  the  general  principles  of  chemistry.  In  fact  one 
of  the  translators  had  just  published  a  more  elementary  text  of 
this  character  which  was  designed  especially  for  first-year  stu- 
dents at  the  Massachusetts  Institute  of  Technology.1 

The  German  text  has,  in  the  main,  been  faithfully  followed, 
although  a  few  minor  changes  have  been  made  to  adapt  the 
book  for  the  use  of  English-speaking  students.  The  German 
authors  have  not  only  cooperated  by  carefully  reading  the 
proof  sheets  but  they  have  also  furnished  the  revisions  and 
additions  which  they  intend  to  incorporate  in  the  Second  Ger- 
man Edition. 

The  book  applies  especially  to  the  more  advanced  college  or 
university  students  who  would  broaden  the  scope  of  their  train- 
ing in  inorganic  chemistry  beyond  that  obtained  from  courses  in 
qualitative  and  quantitative  analysis.  It  might  be  studied  with 
equal  advantage,  simultaneously  with,  or  following,  the  more 
advanced  work  in  analytical  chemistry. 

In  addition  to  the  usefulness  of  this  book  to  students,  it  should 
prove  of  value  to  manufacturing  chemists,  for  although  the 
working  directions  in  the  book  are  given  for  preparations  strictly 
on  the  laboratory  scale,  still  the  direct  bearing  of  the  principles 
and  theories  of  physical  chemistry  upon  the  efficiency  of  the 
chemical  processes  is  brought  out  in  a  manner  which  would 
apply  on  a  large  as  well  as  on  a  small  scale. 

The  translators  desire  to  express  their  deep  obligation  to  Pro- 
fessors Heinrich  and  Wilhelm  Biltz  for  the  interest  they  have 
shown  in  the  translation;  also  to  Mr.  J.  W.  Phelan,  Mr.  C.  B. 
Nickerson,  and  especially  to  Mr.  P.  S.  Fiske  for  their  efficient  aid 
in  reading  and  criticising  the  proof  sheets. 

1  Blanchard:  Synthetic  Inorganic  Chemistry. 
vii 


TABLE   OF   CONTENTS. 


PAGE 

INTRODUCTORY  REMARKS  ON  LABORATORY  PRACTICE 1 


CHAPTER  I. 

The  Elements 9 

REDUCTION  OF"  OXIDES  WITH  CARBON 10 

1.  Lead  from  Lead  Oxide 10 

ALUMINOTHERMY 11 

2.  Manganese  from  Pyrolusite 12 

3.  Chromium  from  Chromic  Oxide 13 

4.  Crystallized  Silicon 13 

5.  "  Crystallized  Boron  " 14 

REDUCTION  WITH  POTASSIUM  CYANIDE 15 

6.  Tin  from  Cassiterite;  Melting-Point  Determination 15 

7.  Pure  Antimony  from  Basic  Antimony  Chloride 17 

REDUCTION  WITH  AQUEOUS  REDUCING  AGENTS 17 

8.  Selenium  Dioxide  and  Pure  Selenium  from  Crude  Selenium ...  17 

9.  Extraction  of  Gold 18 

DESULPHURIZATION  OF  SULPHIDES  BY  THE  PRECIPITATION  PROCESS.  19 

10.  Antimony  from  Stibnite 19 

11.  Mercury  from  Cinnabar;  Sodium  and  Ammonium  Amalgams. .  20 
ROASTING  PROCESSES 20 

12.  Lead  from  Galena 20 

CUPELLATION 21 

13.  Pure  Silver  from  Coin  Metal 21 

METALS  BY  ELECTROLYSIS 22 

14.  Lithium  from  Fused  Lithium  Chloride 22 

CHAPTER  II. 

Changes  of  Condition 26 

POLYMERIZATION  AND  DISSOCIATION 26 

ALLOTROPY 27 

15.  Allotropy  of  Silver  Sulphide 28 

16.  Allotropic  Modifications  of  Sulphur 29 

17.  Transformation  Point  of  Cuprous  Mercuriiodide  and  of  Silver 

Mercuriiodide 31 

ix 


TABLE   OF  CONTENTS. 

PAGE 

THE  PASSIVE  CONDITION 32 

18.  Passive  Iron 32 

AMORPHOUS  STATE 32 

19.  Amorphous  Sulphur 33 

COLLOIDAL  STATE 33 

ADSORPTION  COMPOUNDS 34 

20.  Colloidal  Platinum,  according  to  Bredig 36 

21.  Colloidal  Antimony  Sulphide 37 

22.  Adsorption -of  Iodine  by  Charcoal;  Adsorption  Curve 38 

23.  Lanthanum  Blue 38 

24.  Molybdenum  Blue 39 

25.  Collodial  Gold  Solutions;  Precipitating  Colloids  and  Protec- 

tive Colloids , .  40 

26.  Hydrogels  as  Semipermeable  Membranes 42 


CHAPTER  III. 

Simple  Compounds 44 

OXIDES 49 

27.  Liquid  Sulphur  Dioxide.     Critical  Point 49 

28.  Sulphur  Trioxide  by  the  Contact  Process 51 

Oxidation  of  Naphthalin  with  Sulphuric  Acid 53 

29.  Nitrogen  Dioxide ,  54 

30.  Chromic  Oxide  in  the  Dry  Way  from  a  Chromate 55 

31.  Cuprous  Oxide  from  Fehling's  Solution 56 

HYDRIDES 56 

32.  Cerium  Hydride 56 

33.  Copper  Hydride 57 

ACIDS,  BASES,  AND  SALTS 58 

(a)  Acids  and  Bases 61 

34.  Physico-chemical  Detection  of  Electrolytic  Dissociation 61 

35.  Hydrobromic  Acid 62 

36.  Thallous  and  Thallic  Hydroxides 64 

(b)  Halogen  Compounds 65 

37.  Cupric  and  Cuprous  Bromides 65 

38.  Cuprous  Chloride 66 

39.  Potassium  Iodide 66 

40.  Barium  Chloride  from  Witherite 67 

41.  Manganous  Chloride  from  Waste  Manganese  Liquors ,  68 

42.  Anhydrous  Ferric  Chloride;  Preparation  of  Chlorine 68 

43.  Anhydrous     Ferrous     Chloride;     Preparation     of     Hydrogen 

Chloride. 70 

44.  Anhydrous  Chromium  Trichloride 71 

45.  Sulphur  Chloride,  S2C12 72 

46.  Chlorides  of  Phosphorus 73 

47.  Chlorides  of  Antimony 75 


TABLE  OF  CONTENTS.  xi 

PAGE 

48.  Iodides  of  Bismuth 76 

49.  Bismuth  Tribromide;  Boiling-Point  Determination 77 

50.  Tin  Tetrachloride 79 

51.  Silicon  Tetrachloride 80 

52.  Titanium  Tetrachloride  from  Rutile 81 

53.  Anhydrous  Titanium  Trichloride 84 

(c)  Sulphides 85 

54.  Phosphorus     Pentasulphide;     Thermoelectric     Determination 

of  the  Boiling-Point 85 

55.  Black  Mercuric  Sulphide;  Transformation  into  Cinnabar 88 

56.  Sulphides  of  Tin 88 

57.  Green  Manganese  Sulphide 89 

58.  Titanium  Disulphide 90 

(d)  Nitrides 92 

59.  Hydrogen  Cyanide,  Mercuric  Cyanide,  Cyanogen,  and  Dithio- 

oxamide 92 

60.  Boron  Nitride 94 

61.  Magnesium  Nitride;  Ammonia  from  the  Atmosphere 95 

62.  Chromium  Nitride 96 

(e)  Phosphides 97 

63.  Magnesium  Phosphide 97 

(/)  Carbides. 9" 

64.  Calcium  Carbide;   Acetylene  from  Calcium  Carbide;    Benzene 

from  Acetylene 98 


CHAPTER  IV. 

Compounds  Containing  a  Complex  Negative  Component 99 

COMPOUNDS  WITH  HOMOGENEOUS  COMPLEXES 101 

(a)  Peroxides 101 

65.  Sodium  Peroxide 101 

66.  Barium  Peroxide 102 

67.  Hydrogen  Peroxide 103 

(6)  Polysulphides 106 

68.  Ammonium  Pentasulphide 106 

(c)  Polyhalides 107 

69.  Ammonium  Tribromide 107 

70.  Law  of  Distribution;    Proof  of  the    Existence  of    Potassium 

Tribromide  and  Potassium  Tri-iodide 108 

71.  Rubidium  Iodide  Tetrachloride;    Rubidium  Tri-iodide Ill 

(d)  Polynitrides v Ill 

72.  Sodium  Hydrazoate  from  Sodamide. Ill 

OXYACIDS    AND   THEIR    SALTS 113 

(a)  Cyanates 113 

73.  Potassium  Cyanate;  Urea  from  Ammonium  Cyanate 113 


xii  TABLE  OF  CONTENTS. 

PAGE 

(6)  Oxy-halogen  Acids  and  their  Salts 114 

74.  Electrolytic  Production  of  Sodium  Hypochlorite  and  Potas- 

sium Chlorate 114 

75.  Potassium  Perchlorate 118 

76.  lodic  Acid  and  lodic  Anhydride;  "Time  Reaction" 119 

77.  Potassium  lodate  from  Potassium  Chlorate 120 

78.  Potassium  Bromate  and  Potassium  Bromide 121 

(c)  Nitrites  and  Nitrates 121 

79.  Sodium  Nitrite  from  Sodium  Nitrate 121 

80.  Potassium  Nitrate  from  Sodium  Nitrate 122 

81.  Silver  Nitrate 123 

82.  Bismuth  Nitrate  and  Basic  Bismuth  Nitrate 124 

(d)  Manganates  and  Ferrates 124 

83.  Potassium  Permanganate  by  the  Fusion  Method 124 

84.  Electrolytic  Preparation  of  Potassium  Permanganate 125 

85.  Barium  Ferrate 126 

(e)  Oxyacids  of  Sulphur  and  their  Salts 126 

86.  Sulphuric  Acid  from  Pyrite  by  the  Chamber  Process 128 

87.  Reduction  of  Barium  Sulphate   and  Preparation   of  Barium 

Nitrate 129 

88.  Antimony  Sulphate 129 

89.  Alum  from  Kaolin 130 

90.  Sodium  Thiosulphate 131 

91.  Barium  Dithionate 131 

92.  Sodium  Tetrathionate 132 

93.  Hyposulphurous  Acid 132 

94.  Potassium  Persulphate,  Electrolytically 133 

(/)   Carbonates 134 

95.  Sodium  Carbonate  (Ammonia-Soda  Process) 134 

(g)  Acids  and  Salts  of  Phosphorus 135 

96.  Barium  Hypophosphite 135 

THIOACIDS  AND  THEIR  SALTS 136 

97.  Potassium  Trithiocarbonate  Solution  (Reagent  for  Nickel) 136 

98.  Barium  Trithiocarbonate 137 

99.  Sodium  Thioantimonate 137 

100.  Potassium  Ferric  Sulphide 138 

101.  Ammonium  Copper  Tetrasulphide 139 

COMPLEX  HALOGEN  ACIDS  AND  THEIR  SALTS.     COMPLEX  CYANOGEN 

COMPOUNDS 139 

102.  Hydrofluosilicic  Acid 141 

103.  Potassium  Titanium  Fluoride 142 

104.  Ammonium  Plumbic  Chloride.     Lead  Tetrachloride 143 

105.  Potassium  Lead  Iodide 144 

106.  Potassium  Mercuric  Iodide 144 

107.  Potassium  Cobalticyanide. 145 

108.  Hydroferrocyanic  Acid 146 


TABLE  OF  CONTENTS.  xiii 

PAGE 

109.  Addition   Products   of   Complex    Hydro-metal-cyanic   Acids 

with  Oxygen  Compounds 146 

110.  Cobaltous  Salt  of  Hydromercurithiocyanic  Acid 147 

111.  Potassium  Cobaltothiocyanate 147 

112.  Cadmium  Iodide  (Autocomplex  Compound)  148 

NlTRITO    ACIDS    AND    THEIR   SALTS 149 

113.  Potassium  Mercurinitrite 149 

114.  Sodium  Cobaltinitrite;  Potassium  Cobaltinitrite 149 

115.  Potassium  Tetranitrito-diammine-cobaltate 150 

CONDENSED  ACIDS  AND  THEIR  SALTS 151 

116.  Ammonium  Phosphomolybdate 151 

117.  Silicotungstic  Acid 151 

ORGANOCOMPLEX  COMPOUNDS 152 

118.  Potassium  Ferric  Oxalate;  Platinotypes 152 

119.  Optical  Rotation  of  Uranyl  Lsevo-malate 153 


CHAPTER  V. 

Compounds  Containing  a  Complex  Positive  Component 155 

AMMONIUM  COMPOUNDS  AND  SUBSTITUTED  AMMONIUM  COMPOUNDS  .  156 

120.  Dissociation  of  Ammonium  Chloride 156 

121.  Hydroxylamine    Sulphate;     Hydroxylaminedisulphonate    of 

Potassium;  Acetoneoxime 156 

122.  Hydrazine  Sulphate;  Monochloramine 158 

123.  Semicarbazide  Hydrochloride 161 

124.  Millon's  Base 162 

METAL- AMMONIA  COMPOUNDS 163 

125.  Silver-ammonia  Sulphate 169 

126.  Tetramminecupric  Sulphate  and  Ammonium  Cupric  Sulphate  169 

127.  Tetramminecupric  Chloride 170 

128.  Hexamminenickelous  Bromide 171 

129.  Carbonatotetramminecobaltic  Nitrate 171 

130.  Chloropentamminecobaltic  Chloride  from  the  preceding 172 

131.  Chloropentamminecobaltic  Chloride  from  Cobalt  Carbonate. .  173 

132.  Sulphate  and  Nitrate  of  the  Chloropentammine  Series 174 

133.  Hexamminecobaltic  Salts  (Luteocobalt  Salts) 175 

134.  Aquopentamminecobaltic  Salts  (Roseocobalt  Salts) 176 

135.  Dibromotetramminecobaltic  Bromide  (Dibrompraseo  Salt) ...  178 

136.  1,  2-Dinitritotetramminecobaltic  Salts  (Flavocobalt  Salts) ...  178 

137.  1,  6-Dinitritotetramminecobaltic  Chloride  (Croseocobalt  Salt)  179 

138.  Comparison     of    the    Isomeric    Dinitritotetramminecobaltic 

Salts 180 

139.  Trinitritotriammine  Cobalt 182 

140.  Hexamminechromic  Nitrate  and  Chloropentammmechromic 

Chloride. .  182 


XIV  TABLE   OF  CONTENTS. 

PAGE 

141.  Tetrammineplatinous  Chloroplatinite  (Green  Salt  of  Magnus)  184 

142.  Tetrammineplatinous  Chloride 185 

143.  Isomers  of  Dichlorodiammineplatinum 185 

HYDRATES 186 

144.  Melting-Point  Maximum  for  Magnesium  Nitrate  Hexahydrate ; 

Eutectic  Mixture  of  Barium  Chloride  Dihydrate  and  Water  189 

145.  Calcium  Sulphate  Hemihydrate 193 

146.  Hydrates  of  Sodium  Sulphate;  Supersaturated  Solutions. ...  194 

147.  Transition  Point  of  Sodium  Sulphate 195 

148.  Isomeric  Chromic  Chloride  Hydrates 197 


CHAPTER  VI. 

Complex  Non-Electrolytes 199 

ACID  CHLORIDES 199 

149.  Sulphuric  Acid  Bichloride  and  Sulphuric  Acid  Monochloride  .  199 

150.  Pyrosulphuric  Acid  Chloride 202 

151.  Sulphurous  Acid  Chloride  (Thionyl  Chloride) " 202 

152.  Nitrosylsulphuric  Acid 204 

ESTERS 205 

153.  Ethyl  Nitrate 206 

154.  Amyl  Nitrite  and  Methyl  Nitrite 207 

155.  Symmetrical  Diethyl  Sulphite 207 

156.  Unsymmetrical  Diethyl  Sulphite 209 

157.  Triethyl  Phosphate 210 

158.  Tetraethyl  Silicate 211 

METAL-ORGANIC  COMPOUNDS 211 

159.  Zinc  Ethyl 212 

160.  Lead    Tetraphenyl    by   Grignard's    Reagent;    Diphenyl-lead 

Iodide 213 

161.  Nickel  Carbonyl 215 


CHAPTER  VII. 

Preparation  of  Compounds  of  the  Rare  Elements  from  Their 

Minerals 217 

162.  Lithium  Carbonate  from  Lepidolite,  Petalite  or  Spodumene. 

Spectroscopic  Tests  for  Rubidium  and  Other  Metals 217 

163.  Beryllium  Hydroxide  from  Beryl 219 

164.  Basic  Beryllium  Acetate 220 

165.  Columbium  and  Tantalum  Compounds  from  Columbite 221 

166.  Molybdenum  Compounds  from  Molybdenite 222 

167.  Tungsten  Compounds  from  Wolframite 227 

168.  Working  up  of  Pitchblende  and  Testing  for  Radioactivity 229 


TABLE  OF  CONTENTS.  XV 

PAGE 

169.  Uranium  Compounds 232 

170.  Thorium  Compounds  from  Monazite 234 

171.  Separation  of  the  Rare  Earths 237 

Cerium  Compounds 240 

Lanthanum  Compounds 242 

Didymium  Compounds 244 

INDEX  .  .  245 


LABOKATORY   METHODS 

OF 

INOEGANIC   CHEMISTEY. 


INTRODUCTORY  REMARKS  ON  LABORATORY 
PRACTICE. 

1.   Heating. 

FOR  heating  and  melting  solid  substances  crucibles  made  of 
porcelain,  difficultly-fusible  clay,  or  iron  are  used.  Small 
crucibles  may  be  heated  with  a  Bunsen  burner  or,  for  higher 
temperatures,  with  a  blast  lamp;  for  long  heating  with  the 
latter  it  is  convenient  to  have  a  water  pump  arranged  to 
furnish  an  air  blast.  In  order  to  obtain  a  high  and  uniform 
temperature  small  crucibles  may  be  surrounded  with  a  clay 


Fig.  l. 

mantle;  larger  crucibles  are  placed  inside  a  piece  of  iron  stove- 
pipe or  an  inverted  flowerpot  with  its  hole  enlarged.  Large 
crucibles  may  be  heated  reasonably  hot  with  a  Fletcher  burner 
by  using  such  an  insulating  mantle  to  prevent  some  of  the 
loss  of  heat  by  radiation;  or  to  about  1200°  in  a  gas  furnace. 

1 


INTRODUCTORY. 


If  an  abundant  air  blast  is  available  at  a  fairly  steady  pressure, 
the  most  convenient  and  useful  form  of  furnace  is  the  gas 
furnace,  Fig.  1,  made  of  fire  clay  supported  by  a  casing  of 
sheet  iron,  into  which  the  flame  from  a  large  blast  lamp  is 
admitted  through  a  hole  near  the  bottom.  In  the  absence  of  an 
air  blast,  gasolene  blowpipes,  similar  to  those  commonly  used  by 
plumbers,  serve  almost  as  well  as  the  gas  blast  lamps.  The  use 
of  charcoal  furnaces  is  inexpensive  and  to  be  recommended  even 
at  the  present  day.  Such  a  furnace  is  shown  in  Fig.  2,  one-tenth 
its  natural  size.  It  is  made  of  sheet  iron  and  is  about  250  cm. 
high  and  280  cm.  wide.  Larger  furnaces  may  be  built  of  fire 
brick  and  connected  with  the  chimney  of  the  building.  The 
fire  in  such  a  furnace  can  be  started  with  some  glowing  charcoal 
which  has  been  heated  before  the  blast  lamp;  the  furnace  is  then 
fed  with  layers  of  charcoal  and  coke  and  finally  with  coke  alone. 


Fig.  2. 


Fig.  3. 


For  heating  substances  in  glass  or  porcelain  tubes,  a  flame 
spreader  placed  on  an  ordinary  Bunsen  burner  is  used  when  only 
short  lengths  are  to  be  heated;  for  longer  stretches  a  row  burner, 
as  shown  in  Fig.  3,  is  employed,  and  to  retain  the  heat  a  cover  of 
asbestos  board  can  be  supported  just  above  the  tube  by  means  of 
wires.  The  three  supports  for  the  tube  can  be  protected  by 
wrapping  wet  asbestos  pulp  about  them,  pressing  it  firmly  into 
position  and  allowing  it  to  dry.  For  higher  temperatures  over  a 
considerable  length  of  tube  a  combustion  furnace.,  such  as  is  used 


HEATING. 


3 


in  organic  ultimate  analysis,  is  suitable.  Very  high  temperatures 
may  be  obtained  with  the  blast  lamp  if  the  part  of  the  tube  to 
be  heated  is  surrounded  by  an  asbestos  chamber  like  that  shown 
in  Fig.  4.  The  bottom  of  the  chamber  should  measure  about 
10  X  16  cm.  and  the  height  about  13  cm.  In  the  bottom  is  a 
large  hole  for  the  flame  of  the  blast  lamp  to  enter;  in  the  two 
end  walls  are  holes  through  which  the  tube  to  be  heated  just 
fits.  As  a  cover,  a  piece  of  asbestos  board  serves,  which  is 
loosely  laid  on  top  of  the  box  and  is  provided  near  the  four 
corners  with  round  holes  of  2  cm.  diameter.  The  box  is  held 
together  by  means  of  wire  fastenings.  For  heating  long  tubes 


Fig.  4. 


Fig.  5. 


to  very  high  temperatures,  a  Mitscherlich  tube  furnace  can  be 
used  which  is  50  cm.  long  and  is  made  of  sheet  iron  lined  with 
fire-clay. 

For  providing  and  maintaining  constant  temperatures,  up  to 
1500°,  the  platinum  electrical  resistance  furnace,  and  for  tem- 
peratures above  2000°  the  carbon  resistance  furnace,  are  invalu- 
able. Although  such  furnaces  would  be  very  convenient  for 
some  of  the  preparations  described  in  the  following  pages,  they 
may  be  dispensed  with  on  account  of  their  high  price. 

Many  different  forms  of  apparatus  have  been  devised  for 
uniformly  heating  -flasks  and  similar  vessels  to  moderate  temper- 
atures. Besides  sand  baths,  Babo's  boiling  funnel  (Fig.  5), 
which  is  capable  of  wide  application,  and  nickel  air  baths,  are 
particularly  to  be  recommended. 

When  liquids  are  to  be  boiled  for  a  considerable  length  of 
time,  the  flask  should  be  provided  with  a  condenser  or  a  wide 
glass  tube  held  either  vertically  or  inclined  so  that  vapors  will 
condense  and  the  liquid  run  back  into  the  flask.  This  is  known 
as  heating  with  a  reflux  condenser. 


'4  INTRODUCTORY. 

2.  Evaporation. 

Large  quantities  of  solutions  are  concentrated  in  evaporating 
dishes,  which  give  a  large  surface  for  the  escape  of  vapors  from 
the  liquid.  They  should  be  heated,  unless  otherwise  directed, 
with  a  free  flame  without  wire  gauze  or  asbestos;  but  the  flame 
should  never  be  allowed  to  come  in  contact  with  the  dry  upper 
part  of  the  dish,  as  in  that  case  superheating  would  ensue  which 
might  result  in  the  breaking  of  the  dish  or  in  the  decomposition 
of  the  solid  material  that  separates  on  the  sides.  Thus  as  the 
concentration  progresses  the  flame  must  be  lowered  or  the  evap- 
oration finished  on  the  water  bath.  Small  quantities  of  liquid 
can  also  be  boiled  down  in  beakers;  but  if  this  is  done  it  is 
advisable  to  avoid  bumping  by  stirring  or  by  adding  to  the 
liquid  such  substances  as  small  splinters  of  wood,  bits  of  unglazed 
porcelain,  or  of  pumice. 

3.  Crystallization. 

When  crystals  are  to  be  obtained  from  a  solution,  it  is  often 
necessary  before  the  evaporation  is  finished  to  filter  from  impuri- 
ties. For  this  purpose  an  ordinary  plaited  filter1  is  used 
which  is  placed  in  a  short-stemmed  funnel.  In  case  the  crystals 
separate  rapidly,  a  Biichner  funnel  with  suction  is  employed. 
For  strongly  corrosive  liquids,  hardened  filter  paper  must  be 
used  or,  often  better  still,  a  felt  of  asbestos  fibers;  the  latter  is 
made  by  suspending  the  asbestos  in  water  and  pouring  the  sus- 
pension on  to  the  perforated  bottom  of  a  Biichner  funnel,  or 
Gooch  crucible,  and  drawing  firmly  in  place  by  suction.  Crys- 
tallization will  take  place  on  cooling  if  the  solubility  of  the 
material  decreases  with  sinking  temperature,  otherwise  by 
evaporation  at  a  moderate  heat,  or  at  the  room  temperature  -in 
a  desiccator.  Impure  crusts  tend  to  form  around  the  sides  of 
the  dish  at  the  surface  of  the  liquid;  to  prevent  this  as  much  as 
possible,  dishes  with  perpendicular  walls  (beakers  or  glass  crystal- 
lizing dishes)  should  be  used.  If  this  expedient  is  not  success 
ful,  the  crystals  are  collected  by  themselves  and  the  crusts  dis- 

1  Filters  of  hardened  paper  are  more  expensive  and  allow  the  filtrate  to 
flow  more  slowly  than  those  of  common  paper,  but  they  retain  even  the 
finest  precipitates  and  are  thus  at  times  indispensable. 


DISTILLATION.  5 

solved  again  in  the  mother  liquor.  By  slow  crystallization, 
especially  with  a  deep  body  of  solution,  large  crystals  are 
obtained;  by  more  rapid  cooling,  as  by  shaking  the  solution  in  a 
flask  under  running  cold  water,  a  crystalline  meal  is  obtained 
which,  because  it  contains  less  mother  liquor  inclosed  between 
the  crystal  layers,  is  purer,  although  of  less  characteristic  form. 

4.   Suction. 

Coarse  crystals  can  be  separated  from  the  mother  liquor  by 
sucking  the  mass  through  an  ordinary  funnel  in  which  is  placed 
a  small  glass  marble.  If  a  few  small  crystals  should  run  through 
at  first,  the  liquor  containing  them  is  poured  a  second  time 
through  the  funnel.  Even  fine  crystals  and  powdery  precipi- 
tates can  be  filtered  clear  if  a  thin  cord  of  asbestos  fibers  is  laid 
around  the  marble.  The  suction  funnels  of  R.  Hirsch  and  of 
Biichner  are  also  much  used,  for  all  sizes  of  which  hardened 
filters,  which  do  not,  like  common  filter  paper,  lose  their  fibers, 
are  to  be  obtained  in  the  market. 

If  the  material  on  the  filter  dissolves  but  slowly,  it  can  be 
washed  with  water,  or  with  the  liquid  used  for  the  crystallizing 
medium  if  it  is  other  than  water;  substances  that  dissolve  more 
readily  may  be  washed  with  suitably  diluted  solvents,  for  example, 
with  mixtures  of  alcohol  and  water  and  finally  with  pure  alcohol. 

5.   Distillation.1 

For  distilling  ordinary  liquids,  flasks  are  used;  for  substances 
which  solidify  easily  retorts  are  more  suitable  and  these  in  the  case 
of  high-boiling  materials,  may  be  covered  with  folded  asbestos 
paper  in  order  to  keep  in  the  heat.  The  use  of  a  condenser  is 
usually  superfluous  when  distilling  high-boiling  substances  from 
retorts;  in  distilling  low-boiling  materials  a  condenser  of  suit- 
able length  should  be  joined  to  the  distilling  flask  with  a  cork 
stopper  (Fig.  6),  although  sometimes  the  connection  is  left  open; 
or  a  short  condenser  may  be  slipped  over  the  neck  of  the 
distilling  flask  itself  and  made  tight  with  corks  or  with  pieces 
of  rubber  tubing  (Fig.  7).  In  the  case  of  substances  that  boil 

1  For  fuller  details  regarding  distillation,  especially  of  organic  liquids, 
see  text-books  on  Organic  Chemical  Preparations. 


6 


INTRODUCTORY. 


somewhat  higher,  a  glass  tube  about  the  size  of  one's  finger 
simply  slipped  over  the  side  arm  of  the  distilling  flask  serves  as 
an  air  condenser  (Fig.  8).  If  bad  smelling  or  corrosive  gases 


Fig.  6. 


Fig.  7. 


are  produced  during  the  distillation,  a  suction  bottle  fitted  to 
the  condenser  tube  is  used  as  receiver  and  the  side  arm  of  the 
bottle  is  connected  to  a  tube  which  leads  into  the  ventilating 
flue  of  the  hood. 


DISTILLATION.  7 

If  the  vapors  attack  cork  the  apparatus  can  be  made  tight 
with  wads  of  asbestos  fibers  pressed  into  the  joints,  an  expedient 
which  often  yields  excellent  service. 

For  determining  the  boiling-point  of  the  substance  a  ther- 
mometer should  be  inserted  in  the  neck  of  the  distilling  flask,  or 


Fig.  8. 

in  the  tubulus  of  the  retort,  in  such  a  way  that  its  bulb  will  be 
entirely  surrounded  by  the  vapors  (Figs.  6  and  8);  thus  with  a 
distilling  flask  the  bulb  should  reach  to  the  lower  part  of  the 
neck.  It  is  best  to  use  thermometers  so  short  that  the  mercury 
column  will  remain  entirely  within  the  vapor.  The  thermome- 
ter shows  at  the  beginning  a  temperature  somewhat  too  low 
because  some  time  is  necessary  to  heat  it  to  the  temperature 
of  the  vapor;  toward  the  end  the  reading  becomes  a  little  too 
high  on  account  of  the  vapor  in  the  nearly  empty  flask  becoming 
superheated.  Both  of  these  facts  are  to  be  borne  in  mind  when 
judging  of  the  purity  of  a  substance.  Concerning  fractional 
distillation,  which  is  seldom  of  importance  in  making  inorganic 
preparations,  see  the  preparation  of  acid  chlorides  and  esters, 
Chap.  VI. 

Very  volatile  substances,  such  as  sulphur  dioxide,  can  be  con- 
densed in  common  gas  wash  bottles,  with  ground-glass  joints, 


8  INTRODUCTORY. 

by  means  of  a  freezing  mixture.  When  enough  liquid  has  con- 
densed in  the  wash  bottle  so  that  the  entrance  tube  dips  beneath 
its  surface,  the  further  condensation  takes  place  very  readily.  If 
the  preparation  is  to  be  preserved,  it  is  distilled  into  a  thick-walled 
glass  tube,  of  the  shape  of  a  test  tube,  which  is  drawn  down  at 
its  upper  end  so  that  when  filled  it  can  be  sealed  by  the  blast 
flame  (Fig.  12).  Or  the  substance  can  be  poured  into  a  sealing 
flask  the  neck  of  which  is  of  about  the  thickness  of  a  lead  pencil, 
and  can  be  readily  melted  together. 

6.   Pulverizing. 

Minerals  that  are  to  serve  as  the  raw  materials  for  the  various 
preparations  should,  if  possible,  be  bought  in  powdered  form, 
since  the  crushing  of  large  masses  in  the  laboratory  is  extremely 
laborious.  If  the  pulverizing  must  be  undertaken,  the  material 
should  be  ground  in  small  portions  and  from  time  to  time  the 
fine  powder  should  be  sifted  from  the  coarse  by  means  of  wire 
gauze. 


CHAPTER  I. 


THE   ELEMENTS. 

OCCURRENCE.  The  relative  amounts  of  the  more  common  elements  which 
occur  in  the  earth's  surface,  including  the  oceans  and  the  atmosphere,  is 
shown  by  the  following  table  compiled  from  statistics  by  F.  W.  Clarke: 


Oxygen  

49.98 

Sodium  

2.28 

Phosphorus  

0  09 

Silicon  

25  30 

Potassium  

2.23 

Manganese  

0  07 

Aluminium  

7  26 

Hydrogen  

0.94 

Sulphur  

0  04 

Iron  

5  08 

Titanium  

0.30 

Barium  

0  03 

Calcium  

3.51 

Carbon  

0.21 

Nitrogen  

0  02 

Magnesium  

2.50 

Chlorine  

0.15 

Chromium  

0  01 

It  is  noteworthy  that  certain  well-known  elements,  which  are  important 
from  a  purely  chemical  as  well  as  from  a  technical  standpoint,  are  less 
abundant  than  others  that  have  been  studied  less  and  are  regarded  as 
"rare;"  such  as,  for  example,  titanium,  which  is  widely  distributed  but 
usually  occurs  only  in  small  amounts. 

The  elements,  with  the  exception  of  those  contained  in  the  atmosphere, 
sulphur,  and  the  so-called  "noble"  metals,  are  not  as  a  rule  found  in  a  free 
or  uncombined  state.  The  most  important  natural  compounds  are  the  oxides 
and  sulphides  (e.g.,  of  silicon,  iron,  zinc,  lead,  antimony,  mercury),  the  halides 
(e.g.,  of  sodium,  potassium,  magnesium),  and  the  salts  containing  oxygen 
(silicates,  sulphates  and  carbonates). 

EXTRACTION.  The  free  elements  may  be  separated  from  one  another, 
and  from  impurities,  by  distillation.  In  this  way  the  constituents  of  the 
atmosphere  have  been  isolated,  and  sulphur  is  obtained  from  the  associated 
minerals.  In  other  cases,  the  elements  are  dissolved  out  from  mixtures  by 
means  of  suitable  solvents,  e.g.  gold  and  silver  by  mercury. 

OXIDES  ARE  REDUCED.  The  usual  reducing  agents  are  hydrogen,  car- 
bon or  carbon  monoxide  (No.  1),  aluminium  (Nos.  2  to  5),  magnesium, 
sodium,  potassium  cyanide  (Nos.  6  and  7),  substances  in  solution,  such  as 
sulphurous  acid  (No.  8)  and  finally  the  cathodic  action  of  the  electric  current, 
an  agent  of  very  general  application.  Sulphides  are  either  changed  to  oxides 
by  roasting,  or  they  are  smelted  with  iron  whereby  the  metal  is  formed  in 
the  presence  of  a  slag  consisting  chiefly  of  ferrous  sulphide  (Nos.  10,  11). 
Halides  may  be  decomposed  by  metallic  sodium  or  the  fused  salts  may  be 
electrolyzed  (No.  14). 

9 


10  THE  ELEMENTS. 

From  salts  containing  oxygen,  the  elements  are  usually  obtained  by  an 
indirect  method,  as  for  example  by  first  forming  the  oxides.  Examples 
of  direct  reduction,  however,  are  the  deposition  of  copper  from  copper  sul- 
phate solution  by  means  of  iron,  the  corresponding  precipitation  of  silver 
by  copper,  and  the  preparation  of  phosphorus  by  the  reduction  of  acid  cal- 
cium phosphate  with  carbon  at  a  high  temperature. 

REDUCTION  OF  OXIDES  BY  CARBON. 

Technically,  carbon  is  the  most  important  reducing  agent;  its  first 
product  of  combustion,  carbon  monoxide,  also  has  a  reducing  power  since 
it  is  readily  oxidized  further  to  carbon  dioxide;  in  fact  this  latter  action  is 
often  the  most  important,  as  for  example  in  the  blast-furnace  process  for  the 
reduction  of  iron  from  its  ores.  An  example  of  the  action  of  carbon  mon- 
oxide at  lower  temperatures  is  the  reduction  of  gold  chloride  solutions 
(cf.  No.  25).  The  dissociation  of  carbon  dioxide  at  high  temperatures, 
2  CO2  =  2  CO  +  O2  takes  place  in  opposition  to  the  combustion  of  carbon 
monoxide,  2  CO  +  O2  =  2  CO2.  Accurate  experiments  have  recently  shown 
that  this  dissociation  becomes  appreciable  at  above  1500°  and  increases 
rapidly  with  further  rise  of  temperature,  with  the  result  that  the  combustion 
of  carbon  monoxide,  and  consequently  its  reducing  effect,  becomes  dimin- 
ished. This  contradicts  the  opinion  which  formerly  prevailed  to  the  effect 
that  increase  of  temperature  always  favors  such  technical  reduction  pro- 
cesses. 

1.   Lead  from  Lead  Oxide. 

Place  a  mixture  of  50  g.  litharge  and  3  g.  of  very  fine,  sifted, 
wood-charcoal  in  a  porcelain  crucible,  which  is  from  two-thirds 
to  three-fourths  filled  thereby,  cover  the  mixture  with  3  g. 
of  powdered  borax  glass  and  heat  strongly  over  the  blast  lamp. 
When,  after  about  half  an  hour,  the  reduction  is  complete,  pour 
the  reduced  lead  upon  an  inverted  porcelain  crucible  cover 
which  has  been  previously  heated  so  that  the  hot  lead  will  not 
crack  it.  Yield  40  to  45  g.;  theoretically  46.4  g.  This  process 
is  used  technically  for  recovering  lead  from  litharge  formed  in 
cupellation.1 

Specific  Gravity  Determination.  Fasten  a  clean  piece  of  lead 
weighing  from  5  to  10  g.  in  the  loop  of  a  hair,  or  silk  thread. 
Determine  the  weight  of  the  lead  in  air  (m),  and  the  loss  in 
weight  (w}  when  it  is  entirely  submerged  in  water;  then  if  Q  is 
the  density  of  water  at  the  temperature  of  the  experiment2  and 

1  Cf.  No.  13. 

8  Cf  Tables  of  specific  gravity. 


ALUMINOTHERMY.  11 

y(  =  0.0012  the  density  of  air,  then  d,  the  specific  gravity  of  the 
lead,  is  given  by  the  following  equation: 

d  =  ™  (Q-  A)  +  L 

w 

The  specific  gravity  should  be  computed  to  not  more  than  five 
significant  figures.  The  observations  should  be  checked  with 
the  same  piece  of  lead,  or  with  different  pieces  from  the  same 
preparation.  The  specific  gravity  of  lead  is  about  11.351. 

The  density,  or  specific  gravity,  of  an  element  varies  some- 
what with  the  method  of  preparation.  Thus  with  antimony 
distilled  in  vacuum,  d  =  6.62;  pressed  antimony,  d  =  6.69; 
gold  distilled,  d  =  18.88,  pressed  at  10,000  atmospheres, 

' 


ALUMINOTHERMY. 

Metallic  aluminium  resists  the  action  of  water  and  the  atmosphere,  not 
because  the  metal  is  difficult  to  oxidize,  but  on  account  of  the  fact  that  its 
surface  becomes  covered  with  a  thin,  coherent  layer  of  oxide  which  protects 
it  from  further  attack.  If  the  formation  of  this  layer  of  oxide  upon  the 
aluminium  is  prevented  by  amalgamation,  then  the  aluminium  is  rapidly 
attacked.  Aluminium  vessels  disintegrate  quickly  if  they  are  amalgamated 
with  even  a  trace  of  mercury. 

Remove  the  oil  from  2  to  5  g.  of  aluminium  powder  by  boiling 
it  with  a  little  alcohol;  pour  off  the  latter,  and  cover  the  alu- 
minium with  a  \%  solution  of  mercuric  chloride.  After  a  few 
minutes  decant  off  the  liquid  and  wash  the  powder  several  times 
with  water.  Then  cover  the  amalgamated  metal  with  water  and 
allow  it  to  stand.  Within  a  short  time  there  is  an  evolution 
of  hydrogen,  the  mass  becomes  heated  until  finally  vapors  of 
steam  arise,  and  white  hydrated  aluminium  oxide  is  formed. 
It  is  upon  this  strong  tendency  of  aluminium  to  oxidize  that  the 
processes  of  aluminothermy  are  based. 

Mixtures  of  aluminium  and  oxides  react  together  energetically,  whereby 
the  aluminium  is  converted  into  oxide  and  the  metal  which  originally  was  in 
the  form  of  the  oxide  is  set  free.  Since  one  gram  of  aluminium  on  com- 
bustion yields  over  7000  calories,  almost  as  much  as  carbon,  and  because 
there  are  no  gaseous  products  of  combustion  formed  which  would  carry 
away  heat  from  the  reaction  mixture,  the  temperature  is  raised  to  con- 
siderably above  2000°  C. 

A  mixture  of  ferric  oxide  and  aluminium,  the  "thermite"  of  corrmerce, 


12  THE   ELEMENTS. 

is  used  for  the  rapid  production  of  high  temperatures  in  a  small  space,  thus 
in  welding,  riveting,  etc.,  by  the  Goldschmidt  process.  The  aluminium 
oxide  separates  out  from  the  reaction  in  the  form  of  a  very  hard  crystalline 
substance  which  can  be  used  as  an  abrasive;  often  well-formed  needle-like 
crystals  are  to  be  found  in  the  hollow  spaces  of  the  slag.  Metals  like  man- 
ganese (No.  2),  chromium  (No.  3),  molybdenum  (No.  166)  and  vanadium, 
which  were  formerly  difficult  to  prepare,  or  could  at  best  be  obtained  only  in 
an  unfused  state  and  very  impure,  can  now  be  prepared  without  difficulty  as 
fused  masses,  free  from  carbon,  by  the  reduction  of  their  oxides  with  aluminium. 
Even  silicon  and  boron  can  be  obtained  by  an  alumino-thermic  process, 
if  sulphur  as  well  as  an  excess  of  aluminium  powder  is  added  to  the  oxides. 
The  sulphur  unites  with  aluminium  and  forms  a  slag  of  aluminium  sulphide, 
by  which  reaction  the  high  temperature  required  for  the  reduction  of  the 
oxides  is  reached;  under  the  slag  a  fused  mass  of  metal  is  formed  from  which, 
after  dissolving  away  the  excess  of  aluminium,  crystals  of  silicon,  or  boron, 
are  obtained.  Magnesium,  if  used  in  these  reactions  instead  of  aluminium, 
produces  similar  results. 

2.   Manganese  from  Pyrolusite. 

The  reaction  between  aluminium  powder  and  pyrolusite  is  so 
violent  as  to  be  almost  explosive;  it  is  better  therefore  to  trans- 
form the  pyrolusite  first  into  a  lower  oxide  before  carrying  out  the 
process.  Place  500  g.  of  finely  powdered  and  sifted  pyrolusite 
in  a  Hessian  crucible  and  heat  in  a  charcoal  furnace.  Mix  the 
mangano-manganic  oxide,  Mn304,  thus  obtained  (about  420  g.), 
with  one-third  its  weight  of  aluminium  powder.1  Choose  a 
Hessian  crucible,  of  such  a  size  that  it  will  be  about  three-fourths 
filled  by  the  mixture,  and  embed  it  in  sand  in  a  large,  shallow 
dish.  At  first  add  only  three  or  four  spoonfuls  of  the  mixture 
to  the  crucible;  cover  this  with  5  g.  of  ignition  powder,2  heaping 
it  up  in  the  middle,  and  insert  a  strip  of  magnesium  ribbon  into 
the  powder.  It  is  advisable  for  the  operator  to  wear  colored 
glasses  and  a  heavy  glove,  and  the  experiment  should  be  per- 
formed in  a  place  where  no  danger  can  result  from  flying  sparks. 
When  all  is  ready,  ignite  the  end  of  the  magnesium  ribbon  with 
the  Bunsen  flame;  this  starts  the  reaction.  Then  add  the 
remainder  of  the  charge  from  an  iron  spoon,  not  too  much  at  one 
time,  but  still  rapidly  enough  so  that  the  mass  in  the  crucible  is 
kept  in  a  state  of  brilliant  incandescence  until  the  reaction  is  ended. 

1  Use    the  coarser  aluminium  powder  prepared  especially  for  the  ther- 
mite process,  and  not  the  fine  powder  employed  for  aluminium  paint. 

2  One  part  aluminium  to  10  parts  barium  peroxide. 


CRYSTALLIZED  SILICON.  13. 

After  cooling,  break  the  crucible  and  hammer  away  the  slag 
of  fused  aluminium  oxide  from  the  regulus  of  metallic  man- 
ganese. The  slag  is  very  hard  and  will  scratch  glass.  Yield  of 
manganese,  about  120  g.  Dependent  preparation,  Potassium 
Permanganate,  Electrolytically  No.  84. 

3.    Chromium  from  Chromic  Oxide. 

Chromium  cannot  be  prepared,  at  least  not  on  a  small  scale,  according  to 
the  directions  given  for  manganese,  because  the  heat  liberated  in  the  reaction 
between  chromic  oxide  and  aluminium  is  not  sufficient  of  itself  to  melt  the 
aluminium  oxide.  It  is,  therefore,  in  this  case  necessary  to  heat  the  cruci- 
ble and  its  contents  before  starting  the  reaction,  or  to  add  to  the  mixture  a 
little  chromate  which,  by  reacting  more  violently  with  the  aluminium,  affords 
the  necessary  heat. 

1.  Mix   about  70  g.  of   chromic   oxide  (Ng_,  30)  with   one  or 
two  grams  less  than  the  calculated  amounts  of  aluminium  powder; 
place   the   mixture  in  a  clay   crucible,   which   should  be   about 
two-thirds   filled.      Add    to    the    above  mixture  10  to  15  g.  of 
ignition  powder,  and  place  the  crucible  in  a  glowing  charcoal 
furnace  whereby  the  ignition  powder  ignites,  and  the  mass  enters 
into    violent   reaction.     Finally   remove   the   crucible   from   the 
furnace,  allow  it  to  cool  and  break  it,  together  with  the  slag, 
away  from  the  metallic  chromium.     Yield,  30  to  33  g.     With 
smaller  amounts  of  the  mixture,  a  porcelain  crucible  (which  is, 
however,  very  liable  to  break)  may  be  used  and  heated  over  the 
blast  lamp. 

2.  Treat  an  intimate  mixture  of  70  g.  ignited  chromic  oxide,. 
20  g.  fused  and  powdered  potassium  pyrochromate,  and  32  g.  of 
aluminium  powder,  in  separate  portions  exactly  as  described  in 
the  preparation  of  manganese  (No.  2).     Yield,  30  to  35  g.  of 
chromium  contained  in  one  large  regulus   and  several  smaller 
globules. 

Dependent  preparation,  Anhydrous  Chromic  Chloride,  No.  44. 

4.    Crystallized  Silicon. 

Heat  some  pure,  sifted  sand  in  a  small  evaporating  dish  until 
thoroughly  dried.  Mix  together  90  g.  of  the  dried  sand,  100  g. 
cf  aluminium  powder,  and  120  g.  of  flowers  of  sulphur  in  a  clay 
crucible  which  is  half  filled  thereby.  Cover  the  mixture  with  a 
little  magnesium  powder  and  start  the  reaction  by  igniting  the 


14  THE   ELEMENTS. 

latter.  The  experiment  must  be  performed  out  of  doors,  or 
under  a  hood  with  a  good  draft,  as  considerable  sulphur  diox- 
ide is  evolved.  After  cooling,  break  the  crucible  and  cover  the 
fused  pieces  with  water  in  a  dish  also  placed  under  the  hood; 
hydrogen  sulphide  is  set  free  by  the  hydrolysis  of  aluminium 
sulphide.  The  aluminium  hydroxide  is  easily  rinsed  off  from 
the  regulus  of  metal.  Treat  the  grayish-black,  glistening  metallic 
regulus,  and  any  smaller  globules  of  metal  that  can  be  extracted, 
in  a  beaker  with  strong  hydrochloric  acid,  keeping  it  in  a  warm 
place  for  several  days  and  renewing  the  hydrochloric  acid  from 
time  to  time  until  finally  all  of  the  excess  of  metallic  aluminium 
has  been  dissolved  away  and  a  loose  mass  of  silicon  leaflets 
remains.  Finally  boil  the  crystalline  mass  with  concentrated 
hydrochloric  acid.  Drain  the  crystals  by  suction,  wash  them 
well  with  water,  and  dry  in  the  hot  closet.  A  further  treatment 
of  the  metal  with  hydrochloric  and  hydrofluoric  acids  should  have 
no  effect.  Yield,  20  to  25  g.  Dependent  preparation,  No.  51. 

5.    "Crystallized   Boron." » 

Mix  50  g.  of  anhydrous  boron  trioxide,  75  g.  of  sulphur,  and 
100  g.  of  granulated  or  powdered  aluminium  in  a  crucible  and 
bring  the  charge  into  reaction  in  the  same  manner  as  in  No.  4. 
After  cooling  break  the  crucible  and  the  solidified  melt,  treat  the 
latter  with  water  and  rinse  away  the  aluminium  hydroxide,  formed 
by  hydrolysis,  as  well  as  the  microscopic  crystalline  needles  of 
aluminium  oxide.  Pick  out  the  lumps  of  regulus  and  free  them 
completely  from  the  last  traces  of  slag  both  mechanically  and  by 
long  continued  boiling  with  water.  Treat  the  30  to  40  g.  of 
purified  regulus  particles  thus  obtained  with  concentrated  hydro- 
chloric acid,  adding  only  a  little  at  a  time,  and  after  the  first 
violent  action  is  over  let  stand  in  a  warm  place  several  days 
until  all  of  the  aluminium  has  dissolved  away.  A  heavy,  black, 
shining  mass  of  crystals  is  left.  Remove  the  lighter  impurities 
by  repeated  decantation  with  water;  boil  with  concentrated 
hydrochloric  acid,  placing  a  round  flask  filled  with  cold  water 
over  the  beaker  in  order  to  condense  the  acid  vapors;  then  warm 
with  hydrofluoric  acid  in  a  platinum  dish  for  several  hours,  wash 

1  H.  Blitz,  Ber.  41,  2634  (1908). 


TIN  FROM  CASSITERITE.  15 

again,  and  finally  allow  to  stand  in  a  warm  place  with  dilute 
hydrochloric  acid  until  no  more  bubbles  of  gas  are  given  off. 
The  last  traces  of  aluminium  dissolve  very  slowly. 

About  7.5  g.  of  small,  compact,  mostly  opaque,  black  crystals 
are  obtained  which  have  a  luster  resembling  that  of  hematite. 
The  thinnest  crystals,  which  are  most  often  six-sided,  show  a 
deep  dark-red  color  by  transmitted  light.  The  crystals  scratch 
glass.  They  were  formerly  taken  to  be  pure  crystallized  boron; 
their  composition,  however,  corresponds  to  the  formula  A1B12. 

REDUCTION  WITH  POTASSIUM  CYANIDE. 

6.    Tin  from  Cassiterite;   Melting-point  Determination. 

Native  stannic  oxide,  SnO2,  in  spite  of  the  fact  that  tin  is  closely  related 
to  the  noble  metals,  can  be  reduced  by  carbon  only  at  a  very  high  heat. 
The  reduction  takes  place  far  more  readily  with  potassium  cyanide,  the 
latter  being  oxidized  thereby  to  potassium  cyanate: 

KCN  +  O  =  KCNO. 

This  method,  on  account  of  the  relatively  high  cost  of  potassium  cyanide, 
is  used  only  in  the  laboratory  or  in  the  technical  reduction  of  very  valuable 
metals  from  their  oxides. 

Heat  a  mixture  of  20  g.  very  finely  powdered  cassiterite  and 
20  g.  potassium  cyanide  in  a  porcelain  crucible,  which  should  be 
about  three-quarters  filled  thereby,  for  half  an  hour  over  the 
blast  lamp.  A  clay  mantle  may  be  placed  around  the  crucible 
to  lessen  the  amount  of  heat  lost  by  radiation.  After  cooling, 
wash  the  regulus  of  metallic  tin  (12  to  15  g.)  with  water.  Express 
the  yield  in  per  cent  of  the  weight  of  mineral  taken.  Test  the 
tin  qualitatively  for  iron,  copper,  and  lead,  and  determine  its 
specific  gravity.  The  specific  gravity  of  pure  tin  (in  the  white 

modification)  is  7.287  at  15°. 

i 

Melting-point  Determination. 

The  difference  in  the  energy  content  of  a  substance  in  different  states  of 
aggregation  is  a  still  more  essential  distinction  than  the  mere  outward  char- 
acteristics, —  solid,  liquid,  and  gaseous.  Such  energy  differences  can  be  meas- 
ured in  the  heat  of  fusion  or  of  vaporization. 

The  melting-point  of  a  substance  can  therefore  be  determined  very 
accurately  if  it  is  uniformly  heated  and  the  changes  of  temperature,  as 
shown  by  an  inserted  thermometer,  are  observed.  As  soon  as  the  melting- 


16  THE  ELEMENTS. 

point  is  reached,  the  mercury  remains  stationary  even  although  the  outer 
temperature  is  higher;  the  heat  supplied  to  the  substance  is  all  utilized 
in  causing  its  fusion  before  any  is  available  for  producing  a  further  rise  of 
temperature.  If  the  outer  temperature  is  only  a  few  degrees  higher  than 
the  melting-point,  the  temperature  of  the  substance  remains  constant  for 
a  considerable  time,  so  that  the  thermometer  may  be  read  accurately.  On 
cooling  a  fused  substance,  the  thermometer  likewise  remains  constant  for 
some  time  at  the  solidification  point,  even  when  the  outer  temperature  is 
lower.  With  pure  substances  the  melting-point  and  the  solidification-point 
are  identical. 

Cut  the  tin  into  a  few  small  pieces  and  place  it  in  a  test  tube 
of  1  cm.  diameter  which  dips  4  cm.  into  a  small  beaker  filled 
with  sulphuric  acid.  Cover  the  beaker  with  a  disk  of  asbestos 
board  provided  with  a  hole  into  which  the  test  tube  fits.  Heat 
the  sulphuric  acid  until  the  tin  just  melts;  then  heat  a  ther- 
mometer cautiously  over  a  flame  to  about  200°  and  dip  it  into 
the  molten  metal.  Lower  the  flame  under  the  beaker  so  that 
the  temperature  of  the  bath  falls  slowly,  and,  while  stirring, 
observe  the  changes  of  the  thermometer  reading.  At  the  solidi- 
fication point  the  temperature  remains  constant  for  from  10  to 
20  minutes.  In  determining  this  temperature,  a  supercooling 
occurs  regularly;  i.e.,  the  thermometer  at  first  falls  a  few  degrees 
below  the  solidification  point  while  the  mass  still  remains  per- 
fectly liquid.  Then,  as  the  crystals  of  the  solid  begin  to  sepa- 
rate, the  temperature  rises  rapidly  to  the  solidification  point, 
where  it  remains  perfectly  stationary  while  the  mass  is  crystal- 
lizing. After  this  point  is  determined,  increase  the  flame  a  little 
and  find  the  melting-point  in  an  exactly  corresponding  manner. 
Tin  melts  at  232°  C. 

In  using  a  long,  ordinary  thermometer  the  observed  tempera- 
ture is  always  a  few  degrees  too  low  because  the  entire  thread  of 
mercury  is  not  at  the  desired  temperature.  Such  an  observed 
reading  gives  the  "  uncorrected  melting-point."  It  may  be  cor- 
rected by  adding  the  value,  0.000,16  .a  (t  —  t0),  in  which  a  is 
the  length  in  scale  degrees  of  the  exposed  thread  of  mercury,  t 
is  the  observed  reading,  and  t0  the  average  temperature  of  the 
exposed  mercury.  This  last  value  is  obtained  by  a  second  ther- 
mometer, the  bulb  of  which  is  placed  at  about  the  middle  of  the 
exposed  mercury  column.  The  correction  is  usually  spoken  of 
as  the  "  correction  for  stem  exposure." 


SELENIUM  DIOXIDE  AND   PURE  SELENIUM.  17 

7.   Pure  Antimony  from  Basic  Antimony  Chloride. 

Basic  antimony  chloride  (antimony  oxychloride)  is  a  by- 
product in  the  preparation  of  antimony  trichloride  (No.  47). 
Grind  the  dry  basic  chloride  with  twice  its  weight  of  powdered 
potassium  cyanide,  and  heat  the  mixture  in  a  porcelain  crucible 
over  the  blast  lamp.  To  prevent  loss  of  heat  by  radiation,  sur- 
round the  crucible  with  a  piece  of  iron  stovepipe,  a  flowerpot, 
or  a  larger  graphite  crucible,  the  bottom  of  which  is  cut  off. 
After  cooling  wash  the  metallic  regulus  with  water  until  all  the 
adhering  slag  is  removed.  The  antimony  has  a  silver-white 
appearance  and  shows  a  crystalline  structure. 

Regarding  the  specific  gravity  of  antimony,  see  page  11. 
Dependent  preparation,  No.  88. 

REDUCTION  WITH  AQUEOUS  REDUCING  AGENTS. 

8.   Selenium  Dioxide  and  Pure  Selenium  from  Crude  Selenium. 

From  'a  dilute  aqueous  solution  of  selenious  acid,  the  element  selenium  is 
precipitated  by  soluble  reducing  agents,  such  as  sulphurous  acid  or  hydra- 
zine  salts: 

SeO2  +  2  SO2  =  2  SO3  +  Se. 

A  suitable  solution  of  selenious  acid  may  be  obtained  by  oxidizing  the  raw 
material  with  nitric  acid;  selem'c  acid  is  not  formed  under  these  conditions 
(difference  from  sulphur). 

Place  20  g.  of  crude  selenium  in  an  evaporating  dish,  and 
after  covering  the  dish  with  an  inverted  funnel,  oxidize  by  the 
gradual  addition  of  50  g.  of  concentrated  nitric  acid.  Evap- 
orate the  solution  upon  the  water  bath;  and  after  drying  the 
residue  at  110  to  130°,  sublime  it  from  wide  test-tubes,  or  from 
an  evaporating  dish,  into  a  large  inverted  beaker.  The  subli- 
mate is  usually  somewhat  reddish  in  color,  due  to  the  presence 
of  a  small  amount  of  elementary  selenium;  therefore  treat  it 
once  more  with  nitric  acid,  evaporate  the  solution  to  dryness, 
and  take  up  the  residue  in  water.  Reduce  the  selenious  acid, 
at  the  room  temperature,  by  passing  into  the  solution  sulphur 
dioxide  gas  which  has  been  washed  with  water.  A  red  precipi- 
tate of  selenium  is  slowly  formed;  the  precipitation  takes  place 
more  rapidly,  however,  in  the  presence  of  a  little  hydrochloric 


18  THE  ELEMENTS. 

acid.  The  red,  amorphous  product  can  be  dried  in  vacuo  over 
sulphuric  acid  without  change,  but  by  warming  it  in  contact 
with  the  solution  from  which  it  has  been  precipitated,  it  goes 
over  into  gray  selenium.1 

Selenium  imparts  a  blue  tinge  to  the  Bunsen  flame.  A  porce- 
lain dish  held  in  this  flame  becomes  coated  with  a  brick-red 
(reduction)  spot  which  is  surrounded  by  a  white  (oxidation) 
ring.  At  the  same  time  a  characteristic,  radish-like  odor  is 
noticed,  which  can  be  obtained  even  more  distinctly  by  heating 
selenium  on  charcoal  before  the  blowpipe. 

9.   Extraction  of  Gold. 

Gold  can  be  obtained  mechanically  by  "panning"  the  stamped,  or 
ground  ore,  or,  more  advantageously,  by  treating  the  powdered  ore  with 
mercury  which  dissolves  out  the  gold.  Ores  with  a  small  gold  content,  and 
those  in  which  the  gold  is  chemically  united  with  some  other  element,  are 
treated  either  with  chlorine  water,  whereby  chlorauric  acid  H[AuCl4]  is 
formed,  or  with  potassium  cyanide  solution,  which,  together  with  atmospheric 
oxygen,  converts  the  gold  into  potassium  aurocyanide,  K[Au(CN)J. 

For  the  following  experiments,  employ  a  sand  that  is  very  poor  in  gold ; 
in  case  such  is  not  at  hand,  mix  sand  with  a  little  high-grade  gold  ore,  or 
moisten  the  sand  with  a  gold  solution  and  dry  it  by  ignition. 

Krohncke's  Method  of  Preparing  the  Gold  Solution.  Mix  10 
to  100  g.  of  ore  intimately  with  one-fourth  its  weight  of  sodium 
chloride  (to  form  NaAgCl2  with  the  silver  present)  and  a  little 
potassium  chlorate;  moisten  the  mixture  well  with  concentrated 
hydrochloric  acid  in  a  small  flask,  and  allow  it  to  stand  for  12 
to  24  hours  at  the  room  temperature  with  occasional  shaking. 
Then  heat  it  on  the  water  bath  until  the  greater  part  of  the  free 
chlorine  has  been  expelled.  Finally,  dilute  with  water,  whereby 
any  silver  chloride  is  precipitated,  and  filter.  Evaporate  the 
filtrate  to  a  small  volume  on  the  water  bath,  and  carry  out 
with  this  solution  the  two  reactions  given  below. 

Boring's  Method  of  Preparing  the  Gold  Solution.  Place  the 
sample  of  ore  in  a  glass-stoppered  flask,  add  1  c.c.  of  bromine  and 
about  the  same  amount  of  ether,  and  shake  frequently.  After 
two  hours  —  vapors  of  bromine  must  still  be  visible  —  add 
50  c.c.  of  water  and  allow  the  mixture  to  stand  another  two  hours 


1  Vessels  stained  with  selenium  can  be  cleaned  with  potassium  cyanide 
solution,  whereby  readily  soluble  KCNSe  is  formed. 


ANTIMONY  FROM  STIBNITE.  19 

in  a  warm  place.  Filter  the  solution  and  evaporate  it  to  one- 
fourth  of  its  former  volume. 

Reactions.  To  detect  the  presence  of  gold  in  either  of  the 
above  solutions,  treat  one  part  with  freshly-prepared  ferrous 
sulphate  solution,  whereupon  the  liquid  first  assumes  a  reddish- 
violet  coloration,  and  then  becomes  turbid  through  separation  of 
gold.  With  the  solution  prepared  by  the  Krohncke  method  a 
considerable  quantity  of  the  ferrous  sulphate  should  be  used. 
Treat  a  second  portion  of  the  solution  with  a  few  drops  of  bro- 
mine water  and  a  little  stannous  chloride  solution,  whereby  the 
solution  becomes  colored  at  first  blue,  then  brownish  violet,  and 
later  red;  in  this  test  an  adsorption  compound  of  gold  and 
stannic  hydroxide  is  formed  (Purple  of  Cassius);  cf.  No.  25. 

According  to  the  above  directions,  the  presence  of  about 
0.05  mg.  of  gold  in  the  ore  may  be  detected. 

DESULPHURIZATION  OF  SULPHIDES  BY  THE  PRE- 
CIPITATION PROCESS. 

10.   Antimony  from  Stibnite. 

Sulphides  containing  lead  and  antimony  are  frequently  treated  metallur- 
gically  by  the  so-called  "Precipitation  Process."  By  this  is  understood  the 
fusion  of  the  ore  with  iron  and  suitable  substances  to  form  a  slag,  whereby 
the  iron  serves  as  a  desulphurizing  agent.  The  sulphide  of  iron  formed  dis- 
solves in  the  slag,  and  the  precipitated  metals  collect  at  the  bottom  of  the 
furnace. 

Place  a  mixture  of  100  g.  powdered  stibnite,  42  g.  iron  filings, 
10  g.  anhydrous  sodium  sulphate,  and  2  g.  wood-charcoal  powder 
in  a  Hessian  crucible  and  heat  in  a  charcoal  furnace.  The  tem- 
perature should  not  rise  high  enough  to  melt  the  iron  sulphide 
slag  completely,  but  just  sufficiently  to  soften  it;  this  point  is 
determined  by  stirring  the  fusion  with  an  iron  rod.  After  cool- 
ing and  breaking  the  crucible  a  fused  mass  of  antimony,  weigh- 
ing about  65  g.,  is  found  at  the  bottom.  Test  a  sample  of  this 
crude  antimony  qualitatively  for  the  presence  of  arsenic,  copper, 
iron,  and  lead. 

Purification  I.  Mix  the  finely  powdered  crude  antimony  with 
one-fourth  its  weight  of  powdered  stibnite  and  an  equal  amount 
of  anhydrous  sodium  carbonate,  and  melt  over  the  blast  lamp  in 


20  THE  ELEMENTS. 

a  porcelain  crucible  surrounded  with  a  clay  mantle.     Yield,  60  to 
65  g. 

Purification  II.  To  remove  arsenic  from  the  antimony  puri- 
fied according  to  I,  pulverize  the  metal  again,  mix  it  with  4  g. 
.•sodium  carbonate  and  0.2  g.  potassium  nitrate,  and  melt  the 
mixture  in  the  manner  just  described.  Yield,  50  to  55  g.  of  pure 
antimony.  Dependent  experiment,  No.  88. 

11.   Mercury  from  Cinnabar;  Sodium  and  Ammonium 
Amalgams. 

Heat  a  mixture  of  23  g.  cinnabar  and  somewhat  more  than  the 
calculated  amount  of  iron  filings  in  a  small  retort  of  difficultly- 
fusible  glass.  The  mercury  distils  into  a  small  flask  which  serves 
as  receiver.  Yield,  16  to  18  g. 

Sodium  Amalgam.  Pour  the  mercury  so  obtained  into  a  test 
tube  and  add  clean,  freshly  cut  pieces  of  sodium,  about  the  size 
of  grains  of  wheat,  waiting  each  time  before  adding  a  fresh  piece 
until  the  previous  one  has  reacted;  toward  the  end  assist  the 
reaction  by  heating.  Use  1  g.  of  sodium  in  all.  After  cooling, 
break  the  test  tube  and  collect  the  amalgam,  which  will  keep 
indefinitely  when  preserved  in  well-stoppered  vessels.  A  lump 
of  sodium  amalgam  placed  in  water  gives  a  uniform  evolution  of 
hydrogen. 

Ammonium  Amalgam.  Introduce  a  few  grams  of  sodium 
amalgam  into  20  to  30  c.c.  of  ice-cold,  concentrated  ammonium 
chloride  solution;  the  amalgam  at  once  begins  to  swell,  and  is 
changed  to  a  gray  spongy  mass  of  extremely  voluminous  ammo- 
nium amalgam.  Ammonium  amalgam  decomposes  completely, 
within  a  short  time,  into  hydrogen,  ammonia,  and  mercury. 

ROASTING    PROCESSES. 
12.   Lead  from  Galena. 

By  the  "roasting  process"  galena  is  first  partly  oxidized  at  a  relatively 
low  temperature  to  lead  oxide  or  lead  sulphate;  then  without  any  intro- 
duction of  air  the  mass  is  heated  more  strongly,  whereby  the  oxygen  of  the 
lead  oxide  and  lead  sulphate  accomplishes  the  oxidation  of  all  the  sulphur 
present  to  sulphur  dioxide.  Only  those  ores  are  suitable  for  roasting  which 
contain  but  small  amounts  of  silicates  and  of  sulphides  of  other  metals. 
Silicates  are  harmful,  since  they  give  rise  to  the  formation  of  lead  silicate. 
As  it  is  difficult  to  conduct  the  partial  roasting  satisfactorily  upon  a  small 


PURE    SILVER.  21 

scale,  it  is  better  in  the  following  experiment  to  start  with  a  mixture  of  galena 
and  litharge: 

PbS  +  2  PbO  =  3  Pb  +  SO2, 

PbS  +  PbSO4  =  2  Pb  +  2  SO2. 

Place  an  intimate  mixture  of  20  g.  very  finely  powdered 
and  sifted  galena,  and  37  g.  litharge  in  a  small  clay  crucible, 
and  heat  it  in  a  furnace  as  quickly  as  possible  to  a  bright  red 
heat.  After  about  half  an  hour,  allow  the  contents  of  the 
crucible  to  cool,  and  hammer  away  the  regulus  of  metallic  lead 
from  the  broken  crucible  to  which  it  adheres  because  of  a  glassy 
film  of  lead  silicate.  To  purify  the  lead,  hammer  it  into  a  piece 
as  thick  as  one's  finger,  and  melt  it  in  a  test  tube  while  shaking 
lightly.  Calculate  the  yield  in  per  cent  of  the  theoretical. 

CUPELLATION. 
13.   Pure  Silver  from  Coin  Metal. 

If  an  impure  noble  metal  is  alloyed  with  lead,  and  the  alloy  is  maintained 
in  a  state  of  fusion  in  an  oxidizing  atmosphere,  lead  oxide  is  formed  and 
flows  off,  taking  with  it  the  oxides  of  the  contaminating  metals.  This 
process  is  made  use  of  technically  for  obtaining  pure  silver  from  the  impure 
metal,  and  for  working  up  argentiferous  lead  which  is  obtained  in  the  metal- 
lurgy of  certain  lead  ores;  the  process  is  likewise  used  on  a  small  scale  in  the 
rapid  and  accurate  quantitative  estimation  of  silver  and  gold  in  ores  or  metal- 
lurgical products  (fire-assay). 

Melt  together,  by  means  of  a  slightly  luminous  flame  from  a 
blast  lamp,  about  0.3  g.  of  silver  coin  and  1  g.  of  pure  lead  in  a  flat 
cavity,  made  as  smooth  as  possible,  in  a  piece  of  blowpipe  char- 
coal. Press  some  bone-ash  firmly  into  a  small  porcelain  crucible 
with  a  pestle  so  that  at  the  top  there  is  a  slight  hollow  with  a 
coherent,  perfectly  smooth  surface.  Place  the  metallic  button 
by  means  of  pincers  upon  the  bone-ash  which  is  to  serve  as  a 
cupel.  Support  the  porcelain  crucible  on  a  clay  triangle,  or 
embed  it  in  sand,  with  its  top  slightly  inclined  towards  the  blast 
lamp;  and  direct  the  point  of  the  oxidizing  flame  toward  the 
vicinity  of  the  button,  and  some  of  the  time  directly  upon  it. 
At  first  keep  the  metal  but  barely  melted,  as  otherwise  it  is 
likely  to  spirt.  The  lead  oxide,  as  fast  as  it  is  formed,  runs  off 
and  is  absorbed  by  the  bone-ash;  in  order  that  it  may  not  all 
run  into  the  same  place  change  the  inclination  of  the  crucible 
from  time  to  time.  Continue  the  heating  until  the  size  of  the 


22  THE    ELEMENTS. 

button  no  longer  diminishes;  towards  the  end  apply  a  higher 
heat,  whereby  films  of  metallic  oxide  run  across  the  button  in  the 
direction  of  the  oxidizing  draft,  and  lead  oxide  separates  out  on 
the  opposite  side  in  the  form  of  dark-brown  crystals  containing 
copper.  The  end  of  the  process  is  reached  upon  the  disappear- 
ance of  the  oxide  film,  which  at  the  last  shows  for  an  instant  a 
rainbow-like  play  of  colors.  This  is  known  as  the  "  blick."  After 
this  point  is  reached,  further  heating  causes  a  loss  of  silver  by 
volatilization.  During  the  process  a  part  of  the  lead  is  volatilized. 
The  metal  still  contains  a  small  amount  of  copper.  Melt  it 
again  with  1  g.  lead,  and  cupel  the  greater  part  of  this  lead  away 
as  before,  using  a  fresh  cupel.  Free  the  button  from  bone-ash 
by  hammering  it  into  a  cube  (holding  it  in  pincers),  and  remove 
the  last  traces  of  lead  by  heating  it  on  a  third  cupel.  Weigh  the 
silver  button  that  is  finally  obtained.  Yield,  96  to  97%  of  the 
actual  silver  content  of  the  coin. 

METALS    BY    ELECTROLYSIS. 
14.   Lithium  from  Fused  Lithium  Chloride. 

Fused  lithium  chloride,  like  the  aqueous  solution  of  the  salt,  is  largely 
dissociated  electrolytically  ;  and  if  it  is  subjected  to  the  action  of  an  electric 
current,  the  free  elements,  lithium  and  chlorine,  separate  at  the  electrodes  in 
accordance  with  the  following  equations,  in  which  ©  and  ©  represent 
definite  amounts  of  electricity: 

Li  ©        +  Q  Li       +  ©©(reaction  at  the  cathode) 

lithium  negative  neutral 

ion  electricity  electricity 

Cl  0        +  ©  Cl       +  ©  ©   (reaction  at  the  anode) 

chlorine  positive  neutral 

ion  electricity  electricity 

According  to  Faraday's  Law,  equal  amounts  of  electricity  cause  the 
deposition  of  equivalent  amounts  of  different  substances;  and,  in  fact,  as 
measurements  have  shown,  96,540  ampere-seconds  l  discharge  one  gram- 
equivalent  of  a  substance.  This  amount  of  electricity,  therefore,  will  under 
the  most  favorable  conditions  (i.e.,  when  all  of  the  current  passing  through 
the  fused  lithium  chloride  is  utilized  for  the  purpose  of  electrolytic  deposition) 
liberate  one  atomic  weight  in  grams  of  both  lithium  and  chlorine.  Thus, 
for  example,  if  a  current  of  1  ampere  should  pass  for  one  hour,  the  quantity 

,  ....  .  .  ,  ,     7.03x3600       noco 

of  lithium  deposited  would  be  —  T^-TTT-    =  0.262  g. 

~ 


One  ampere-second  =  one  coulomb;  96,540  coulombs  =  one  Faraday. 


LITHIUM   BY  ELECTROLYSIS. 


23 


In  carrying  out  electrolytic  preparations,  the  apparatus  is  usu- 
ally arranged  in  accordance  with  the  following  scheme  (Fig.  9), 
which  is,  in  principle,  identical  with  that  employed  in  electro- 
analysis.  The  decomposition  cell,  an  ammeter,  and  a  regulating- 
resistance  are  introduced  in  series  into  an  electric  circuit.  The 
binding  clamps  of  the  cell  may  also  be  connected  with  wires  from 
a  voltmeter  in  order  to  determine  the  difference  in  potential 
between  the  two  poles  during  the  electrolysis.  The  ammeter 
should  remain  in  the  circuit  throughout  the  whole  of  the  elec- 
trolysis; the  voltage  is  measured  only  from  time  to  time,  and 
has  less  value  for  determining  the  course  of  the  reaction 
when  one  is  working  according  to  well-tested  directions.  A 
storage  battery  may  serve  as  the  source  of  electricity;  for  meas- 
uring the  current  and  voltage  the  simpler  and  inexpensive  instru- 
ments suffice,  but  they  should  be  compared  with  instruments  of 
precision. 


Battery 


Resist 


Decomposition  Cell 


FIG.  9. 


In  all  cases  where  a  metal  which  decomposes  water  is  to  be 
prepared,  a  fused  compound  rather  than  its  aqueous  solution  must 
be  used  for  the  electrolyte.  The  melting-point  of  the  salt  may  be 
lowered  by  admixture  with  some  other  suitable  salt. 

Mix  together  30  g.  of  dry  lithium  chloride  and  an  equal  weight 
of  dry  potassium  chloride  in  a  porcelain  crucible,  6  to  8  cm. 
high  and  9  cm.  in  diameter.  Melt  the  mixture  and  then  after 


24  THE   ELEMENTS. 

starting  the  electrolysis  maintain  a  sufficient  flame  to  supplement 
the  heating  effect  of  the  current  in  keeping  the  mass  just  liquid. 
For  the  anode,  use  a  rod,  0.8  cm.  in  diameter,  of  arc-lamp  car- 
bon which  is  not  attacked  by  the  liberated  chlorine;  and  for  the 
cathode  use  an  iron  rod  0.3  cm.  thick.  Insert  the  cathode 
through  a  cork  in  the  upper  end  of  a  glass  tube,  of  2  cm. 
diameter,  and  place  the  whole  in  the  fused  electrolyte  so  that 
the  glass  tube  dips  1  cm.  below  the  surface  and  the  iron  rod 
0.3  cm.  deeper  (Fig.  9).  The  metallic  lithium  collects  in  the 
space  between  the  iron  wire  and  the  glass  tube,  and  the  glass 
mantle  protects  it  from  being  disseminated  throughout  the  elec- 
trolyte. When  all  is  ready  close  the  circuit,  noting  the  time,  and 
regulate  the  resistance  so  that  a  current  of  between  6  and  10 
amperes  passes.  In  order  to  obtain  this  current,  the  voltage  of 
the  storage  battery,  since  it  has  the  resistance  of  the  electrolyte 
to  overcome,  must  not  be  too  small;  it  should  be  between  7  and 
12  volts  (3  to  6  accumulator  cells  connected  in  series).  During 
the  electrolysis  note  frequently  the  time  and  amperage,  and  from 
these  readings  compute  the  number  of  ampere-seconds  by  multi- 
plying the  average  reading  of  the  ammeter  by  the  number  of 
seconds  which  have  elapsed. 

When  the  reaction  is  progressing  with  moderate  strength,  but 
without  being  disturbed  by  too  vigorous  an  evolution  of  chlorine, 
metallic  lithium  can  be  observed  collecting  at  the  cathode 
in  the  space  between  the  iron  rod  and  the  glass  tube.  From 
time  to  time  flashes  of  light  occur  at  the  anode  which  are  prob- 
ably due  to  an  insulating  envelope  of  chlorine  gas  which  is 
formed  about  the  carbon  and  causes  a  marked  lessening  of  the 
current;  this  difficulty  is  easily  remedied  by  occasionally  break- 
ing the  circuit  for  a  moment.  At  the  end  of  twenty  minutes  a 
considerable  amount  of  lithium  should  have  collected.  Raise 
the  cathode  from  the  fused  salt,  holding  an  iron  spoon  to  pre- 
vent any  lithium  from  falling  out  of  the  glass  tube,  and  dip  the 
whole  under  petroleum.  After  cooling  remove  the  lithium  from 
the  tube  with  a  knife  and  weigh  it  under  petroleum.  The  current 
yield,  i.e.,  the  yield  reckoned  on  the  amount  which  the  current 
should  theoretically  produce,  is  about  70%. 

The  lithium  thus  prepared  is  contaminated  with  a  little  potas- 
sium. Determine  the  specific  gravity  by  means  of  a  pycno meter, 


LITHIUM  BY  ELECTROLYSIS.  25 

using  petroleum  of  known  density.  For  this  purpose  weigh  the 
pycnometer  (1)  empty,  (2)  filled  with  petroleum,  (3)  containing 
the  lithium  alone,  and  (4)  containing  the  lithium  and  enough 
petroleum  to  fill  the  instrument  completely.  From  these  weigh- 
ings and  the  known  specific  gravity  of  the  petroleum,  the  density 
of  the  lithium  can  be  computed;  it  should  be  0.534. 


CHAPTER  II. 


CHANGES   OF   CONDITION. 

CHANGES  in  the  state  of  aggregation  of  substances  have  already  been  men- 
tioned in  the  preceding  chapter;  but  under  changes  in  condition  are  included 
also  polymerization  and  dissociation,  as  well  as  the  formation  of  allotropic, 
passive,  amorphous,  and  colloidal  modifications.  These  phenomena  are,  with 
few  exceptions,  not  confined  to  any  particular  class  of  substances. 

POLYMERIZATION  AND  DISSOCIATION. 

The  phenomena  of  polymerization  and  dissociation,  as  far  as  they 
occur  with  gaseous  or  dissolved  bodies,  are  capable  of  a  complete  theoret- 
ical interpretation.  By  polymerization  is  understood  the  adding  together  of 
particles  of  material  of  the  same  kind  to  form  larger  aggregates;  the  opposite 
of  this  process  is  dissociation.  The  occurrence  of  either  may  be  demon- 
strated by  determining  the  molecular  weight  of  the  substance  in  the  vaporized 
or  dissolved  state.  For  example,  the  diatomic  iodine  molecules  dissociate 
at  high  temperatures  into  atoms;  on  cooling,  the  atoms  again  poly- 
merize into  molecules  of  the  original  sort: 

Dissociation 

I,  =  21 
Polymerization 

In  a  similar  way  ferric  chloride  forms  simple  molecules  at  high  temperatures, 
and  double  molecules  at  lower  temperatures: 

Dissociation 
Fe2Cl6  =  2  FeCl3 

Polymerization 

The  enumeration  of  experiments  illustrating  these  phenomena  may  be 
dispensed  with  at  this  point;  the  conception  of  dissociation,  however,  is 
developed  more  fully  in  the  general  section  of  Chapter  III,  and  that  of  elec- 
trolytic dissociation  in  the  sections  on  acids,  bases,  and  salts,  and  in  No.  34. 


26 


ALLOTROPY.  27 


ALLOTROPY.1 

Many  solid  substances  appear,  under  different  conditions,  in  two  or  more 
•distinct  forms  (known  as  allotropic  modifications),  which  are  distinguished 
from  one  another  by  color,  density,  crystalline  form,  solubility,  and  other 
physical  properties.  The  best  known  instances  of  this  are  shown  by  carbon 
and  phosphorus. 

In  the  same  way  that  different  states  of  aggregation  are  separated  by  a 
temperature  boundary  which  is  dependent  on  the  pressure,  so  also  for  the 
mutual  transformation  of  numerous  allotropic  forms,  there  is  a  definite  tem- 
perature, dependent  only  on  the  pressure,  below  which  the  one,  above  which 
the  other  form  is  stable.  This  is  known  as  the  transition  temperature. 

In  addition  to  this  sort  of  allotropy,  which,  because  the  two  forms  can 
change  simultaneously  into  one  another  and  can  exist  in  equilibrium  at  the 
transition  temperature,  is  known  as  enantiotropy,  there  is  also  a  second  sort 
known  as  monotropy.  Two  modifications  of  a  substance  are  monotropic 
when  the  one  changes  into  the  other,  but  the  latter  cannot  change  back 
directly  into  the  first.  They  possess  no  transition  temperature;  one  form  is 
under  all  conditions  less  stable  than  the  other,  and  therefore  is  only  to  be 
observed  in  virtue  of  the  extreme  slowness  with  which  the  transformation 
takes  place. 

An  example  of  enantiotropy  is  shown  by  sulphur,  which  above  96°  is 
monoclinic,  below  96°  rhombic.  An  example  of  monotropy  is  furnished  by 
iodine  chloride,  IC1,  the  labile  form  of  which  melts  at  14°  and  the  stable 
form  at  27°.  (Cf.  No.  55.) 

The  energy  difference  between  allotropic  modifications  corresponds  entirely 
with  that  existing  between  the  solid  and  the  liquid  states  and,  like  the  latter, 
is  measured  by  the  heat  of  transformation. 

For  determining  the  transition  point  of  allotropic  forms,  the  following 
methods  are  chiefly  used: 

(1)  The  THERMIC  METHOD.     The  point  on  the  heating  or  cooling  curve  is 
•determined  at  which,  in  consequence  of  the  absorption  or  setting  free  of  the 
heat  of  transition,  a  retardation  in  the  otherwise  steady  rise  or  fall  of  tem- 
perature occurs;    compare  the  determination  of  the  melting-point  of    tin, 
-No.  6,  also  Nos.  15  and  16. 

(2)  The  DILATOMETRIC  METHOD,  which  depends  on  a  comparison  of  the 
•densities  (or  of  the  volumes)  on  both  sides  of  the  transition  point. 

(3)  The  OPTICAL  METHODS,  which  frequently  permit  a  very  sharp  observa- 
tion of  a  change  irf  crystalline  form,  or  in  color,  at  the  transition  point. 
(No.  17.) 

(4)  ELECTRICAL  METHODS,     (a)  The  electrical  conductivity  of  a  substance 
is  plotted  graphically  against  the  temperature  and  the  point  of  inflection  of 
the  curve  determined. 


1  Concerning  allotropy,    see   also   B.   Roozeboom,   Heterogene  Gleichge- 
wichte,  Vol.  1,  p.  109  (1901). 


28 


CHANGES   OF  CONDITION. 


(b)  The  temperature  is  found  at  which  the  difference  of  potential  between 
each  of  the  two  modifications  and  some  common  electrolyte  is  the  same;  in 
other  words,  the  electromotive  force  of  an  element  constructed  from  the  twa 
modifications  and  a  common  electrolyte  is  equal  to  zero. 

15.    Allotropy  of  Silver  Sulphide. 

Silver  sulphide  has  a  transition  temperature  of  above  170°,  which  can  be 
determined  both  from  the  electrical  resistance  and  from  the  cooling  curve. 
The  phenomenon  of  supercooling  not  being  very  pronounced  with  this  sub- 
stance, a  point  of  inflection  is  found  on  the  cooling  as  well  as  on  the  heating 
curve;  on  the  former  it  lies  naturally  at  a  lower  temperature  than  on  the 
Utter. 

In  order  to  prepare  silver  sulphide,  treat  a  hot  solution  of  about 
20  g.  of  silver  nitrate  in  300  to  400  c.c.  of  water  with  hydro- 
gen sulphide,  and  wash  the  precipitate  by  decantation.  After 


210 

no 
no 

210° 

200 
190 
ISO 
170 
100 


130". 


220 
210 
200° 
190° 
180° 
170° 
160° 
150° 
140° 
130C 


100  200  300  400   500 
Seconds 

FIG.  10. 


100   200  300  400  600   COO 
Seconds 

FIG.  11. 


removing  the  liquid  by  suction,  dry  the  preparation  in  the  hot 
closet.  Silver  sulphide  thus  formed  is  not  absolutely  free  from 
uncombined  sulphur. 

Place  about  8  g.  of  the  silver  sulphide  in  a  test  tube  and  insert  a 
thermometer.  Heat  the  test  tube  to  about  290°  in  an  air  bath, 
consisting  of  a  porcelain  crucible  covered  with  a  piece  of  asbestos 
board.  Some  sulphur  sublimes,  but  this  does  not  interfere  with 
the  experiment.  Then  transfer  the  test  tube  quickly  to  a  small 
beaker  containing  sulphuric  acid  or  paraffin  at  145°,  which  tem- 
perature must  be  kept  constant  to  within  a  few  degrees  throughout 


ALLOTROPIC   MODIFICATIONS   OF   SULPHUR.  29 

the  experiment.     Read  the  temperature  every   10  seconds  and 
plot  the  results  (Fig.  10). 

When  the  temperature  has  fallen  to  about  130°,  proceed  to 
obtain  the  heating  curve  by  warming  the  bath  so  that  its  tem- 
perature keeps  constantly  20°  or  30°  above  that  of  the  sulphide 
(Fig.  11);  the  continuous  curve  gives  the  bath  temperatures,  the 
dotted  one  the  readings  for  the  silver  sulphide.  The  point  at 
which  the  transformation  is  finished,  and  the  temperature  begins 
to  rise  rapidly  is  particularly  sharp. 

16.    Allotropic  Modifications  of  Sulphur. 

Monoclinic  sulphur  has  a  density  of  1.96  and  a  melting-point  of  119.25°; 
rhombic  sulphur  a  density  of  2.06  and  a  melting-point  of  112.8°.  When  sul- 
phur crystallizes  from  a  solution  in  carbon  bisulphide,  it  is  obtained  in  the 
rhombic  modification,  which  at  room  temperature  can  be  preserved  unchanged 
for  an  indefinite  time.  If  on  the  other  hand  melted  sulphur  is  allowed  to 
solidify,  the  crystals  which  form  are  not  of  the  rhombic,  but  are  commonly  of 
the  monoclinic,  modification;  this  modification  persists  for  some  time,  even 
below  the  transition  temperature,  as  an  unstable  form,  for  the  reason  that  the 
transition  takes  place  slowly.  In  addition  there  have  been  six  or  seven  other 
allotropic  forms  of  sulphur  shown  to  exist  by  crystallographic  optical  methods. 

Monoclinic  Sulphur.  Transformation  into  Rhombic  Sulphur. 
Heat  about  7  g.  of  sulphur  in  a  test  tube  until  it  has  melted  and 
begun  to  turn  dark-colored  (140°  to  150°);  dip  a  thermometer 
in  the  liquid  and  clamp  it  in  position.  Allow  the  melted  sulphur 
to  cool  in  a  sulphuric-acid  bath  at  80°  to  90°,  and  when  it  is  at 
110°  arrest  the  supercooling  by  dipping  a  glass  thread  into  the 
melt.  The  sulphur  crystallizes  to  a  wax-yellow  mass  of  trans- 
parent monoclinic  needles,  while  the  liberated  heat  of  solidifica- 
tion raises  the  temperature  several  degrees.  By  the  next  day, 
or  more  quickly  on  moistening  the  mass  with  carbon  disulphide, 
the  crystals  become  light-yellow  and  opaque,  changing  thereby 
into  an  aggregate  of  rhombic  crystals. 

Transformation  of  Rhombic  into  Monoclinic  Sulphur.  Heat  a 
few  clear  crystals  of  rhombic  sulphur  for  two  or  three  hours  in  a 
test  tube  which  dips  in  a  bath  of  a  boiling,  concentrated  solution 
of  common  salt  (temperature  108°  to  112°),  replacing  when  neces- 
sary the  water  evaporated  from  the  bath.  The  sulphur  crystals 
gradually  become  clouded,  and  change  finally  to  a  friable,  light- 
yellow  mass  of  monoclinic  sulphur. 


30  CHANGES  OF  CONDITION. 

The  Transition  Temperature  from  the  Heating  Curve.  The  cool- 
ing curve  of  sulphur  does  not,  when  determined  in  the  simplest 
manner,  show  an  exact  point  of  solidification,  because  this  is 
dependent  on  the  temperature  to  which  the  material  has  pre- 
viously been  heated;  nor  does  it  indicate  sharply  the  transition 
temperature,  on  account  of  the  ability  already  mentioned  of 
rnonoclinic  sulphur  to  persist  in  the  unstable  condition.  Both 
values,  however,  can  be  established  within  5°  if  the  rate  is 
measured  at  which  the  temperature  of  rhombic  sulphur  rises  with 
a  uniform  application  of  heat. 

Melt  about  7  g.  of  sulphur  in  a  test  tube,  place  the  bulb  of 
a  thermometer  in  the  liquid,  let  the  latter  solidify,  and  after 
moistening  it  with  a  little  carbon  bisulphide  allow  it  to  stand  until 
the  next  day.  It  is  well  also  to  prepare  one  or  two  duplicate 
tubes  in  the  same  manner.  Heat  one  of  the  tubes  in  a  sulphuric- 
acid  bath  in  such  a  manner  that  the  temperature  of  the  bath 
keeps  constantly  about  15°  in  advance  of  that  shown  by  the 
thermometer  in  the  tube,  and  never  rises  more  rapidly  than  1°  in 
10  seconds.  Read  the  thermometer  every  10  seconds  between 
60°  and  120°,  and  plot  the  corresponding  temperatures  and  times 
on  coordinate  paper.  The  rise  in  temperature  is  retarded 
between  95°  and  100°  (the  transition  point)  in  consequence  of 
the  amount  of  heat  absorbed  in  the  transformation.  In  the 
further  course  of  the  curve  there  is  only  a  moderately  rapid  rise 
in  temperature,  —  partly  because  the  transition  is  not  completed, 
partly  because  melting  begins  —  until,  above  a  point  between 
115°  and  120°  (the  melting-point),  the  temperature  again  rises 
rapidly. 

Allow  the  melt  to  cool  again,  and  induce  crystallization 
between  110°  and  100°  by  means  of  a  glass  thread,  or  by  seeding 
with  a  minute  monoclinic  crystal;  then  allow  to  cool  further  to 
60°,  and  repeat  the  experiment.  This  time,  since  only  mono- 
clinic  sulphur  is  present,  there  is  no  indication  of  a  transition 
point;  the  melting-point,  however,  is  shown,  although,  as  in  the 
preceding  experiment,  it  is  not  sharply  defined. 

According  to  recent  investigations,  it  seems  possible  that  a  relation  may 
be  shown  between  the  characteristic  changes  in  consistency  and  color  which 
sulphur  undergoes  on  further  heating  and  the  existence  of  allotropic  liquid 
modifications.  This  would  constitute,  if  we  disregard  the  so-called  "liquid 


TRANSITION   POINT.  31 

crystals,"  the  first  example  of  allotropy  in  the  liquid  state.  The  views  held 
concerning  these  phenomena  are,  however,  at  the  present  time  very  con- 
tradictory. 

When  sulphur  is  heated  in  a  test  tube  above  160°  it  becomes 
dark  and  so  viscous  that  the  tube  can  be  inverted  without  the 
sulphur  flowing  out. 

17.   Transition  Point  of  Cuprous  Mercuriiodide  Cu2  [HglJ  and  of 
Silver  Mercuriiodide  Ag2  [HglJ  . 

First  prepare  mercuric  iodide  by  precipitating  a  solution  of 
6.8  g.  of  mercuric  chloride  with  a  solution  of  8.3  g.  of  potassium 
iodide.  Wash  the  precipitate  once  by  decantation,  and  dissolve  it 
together  with  8.3  g.  of  potassium  iodide  in  50  c.c.  of  water.  Mix 
the  filtered  solution  with  another  filtered,  concentrated  solution 
containing  12  g.  of  blue  vitriol,  and  pass  sulphur  dioxide  into  the 
filtrate,  in  order  to  reduce  the  cupric  salt.  The  sulphur  dioxide 
can  be  prepared  from  sodium  sulphite  and  sulphuric  acid  (cf.  foot- 
note, page  71).  A  bright-red  precipitate  is  produced  of  some- 
what the  appearance  of  mercuric  iodide.  Wash  it  thoroughly  on 
the  suction  filter  and  dry  it  in  the  hot  closet.  Yield,  about  20  g. 

Cuprous  mercuriiodide  is  transformed  at  about  71°  into  a 
black  modification.  The  color  change  can  be  observed  if  a  pinch 
of  the  material  is  gently  heated  in  a  test  tube  over  a  free  flame; 
on  cooling,  the  black  color  changes  again  to  red.  To  determine 
the  transition  temperature,  a  sample  may  be  heated  slowly  in  a 
dry  test  tube  which  'is  immersed  in  a  beaker  of  water  containing 
a  thermometer  that  serves  also  as  a  stirrer.  By  means  of  several 
repetitions,  the  transition  temperature  may  be  obtained  in  this 
way  with  considerable  accuracy. 

The  silver  salt  also  of  hydromercuriiodic  acid,  H2[Hg  I J,  pos- 
sesses a  transition  point  between  40°  and  50°,  which  is  likewise 
characterized  by  a  change  in  color.  Below  this  temperature  the 
salt  is  yellow;  above,  it  is  red. 

Precipitate  a  few  cubic  centimeters  of  mercuric  chloride  solu- 
tion with  a  solution  of  potassium  iodide,  and  redissolve  the  pre- 
cipitate in  an  excess  of  the  precipitant.  Then  add  a  few  drops  of 
silver  nitrate  solution  and  observe,  without  filtering,  the  change  in 
color  of  the  precipitate  on  heating  and  cooling.  This  transition, 
point  is  less  sharp  than  that  of  the  copper  salt. 


32  CHANGES    OF    CONDITION. 


THE  PASSIVE   CONDITION. 

The1  change  of  some  metals,  particularly  chromium  and  iron,  into  the 
passive  condition  is  a  phenomenon  distinct  from  allotropy,  and  one  of  which 
the  significance  is  not  yet  fully  understood.  The  surface  of  metals  when 
passive  is  more  "noble"  than  when  they  exist  in  the  ordinary  state,  as  is 
shown  by  the  lesser  tendency  to  react  chemically,  and  by  the  electrochemical 
behavior;  in  other  respects  there  is  no  distinction  between  the  ordinary  and 
the  passive  conditions. 

Metals  can  be  obtained  in  the  passive  state  by  treatment  with  nitric  acid, 
or  other  oxidizing  agents,  or  by  anodic  oxidation. 


18.   Passive  Iron. 

Fill  a  small  beaker  with  concentrated  nitric  acid,  a  large  one 
with  water,  and  another  small  one  with  copper  sulphate  solution. 
Suspend  a  piece  of  iron  4  to  5  cm.  long  (a  thick  iron  screw  will 
answer)  bv  a  platinum  wire  and  submerge  it  completely  in  the 
nitric  acid,  then  lift  it  cautiously  from  the  acid,  rinse  it  in  the 
water,  and  finally  dip  it  into  the  copper  sulphate  solution,  using 
care  in  each  operation  not  to  allow  the  iron  to  come  in  contact 
with  any  solid  object,  nor  to  suffer  any  jar.  When  the  iron  is 
taken  from  the  copper  sulphate  solution  it  shows  a  gray  color; 
it  has  not  reacted  with  the  copper  salt.  The  reaction  takes  place 
immediately,  however,  when  the  passive  condition  is  destroyed; 
for  example,  by  striking  a  sharp  blow  with  a  glass  rod.  Starting 
from  the  point  hit,  a  coating  of  copper,  which  precipitates  from 
the  adhering -film  of  copper  salt  solution,  spreads  over  the  sur- 
face of  the  iron. 

The  experiment  can  be  repeated  immediately,  for  the  coating 
of  copper  is  quickly  dissolved  in  the  nitric  acid.  When  carried 
out  on  a  larger  scale  this  experiment  is  well  adapted  for  lecture 
demonstration. 

AMORPHOUS    STATE. 

As  already  stated,  the  phenomenon  of  supercooling  is  very  generally  observed 
when  melted  substances  solidify.  This  can  often  be  obviated  by  means  of 
"seeding"  with  small  fragments  of  the  crystallized  substance.  In  other 
cases  the  supercooling  persists  far  below  the  true  temperature  of  crystalli- 
zation, and  a  gradual  changing  into  the  crystalline  form  takes  place  only  very 
slowly,  or  perhaps  not  even  to  a  detectable  extent.  Such  strongly  super- 
cooled fluids  are  designated  as  amorphous  substances  when  they,  like  glass 


AMORPHOUS  SULPHUR.  33 

or  obsidian,  possess  such  a  viscosity  that  they  appear  to  be  solid  bodies. 
The  amorphous  solid  condition  is,  therefore,  not  distinguished  from  the 
fluid  state  by  any  discontinuity,  but  is  so  distinguished  from  the  solid 
crystalline  condition.  It  is  possible  by  certain  expedients,  for  example  by 
continued  heating  at  just  below  the  melting-point,  to  hasten  the  rate  of 
crystallization,  as  in  the  devitrification  of  glass.  With  substances  which  are 
capable  of  solidifying  in  either  a  crystalline  or  amorphous  condition,  the 
quantity  of  crystals  obtained  is  for  this  reason  greater,  the  more  slowly  the 
cooling  takes  place. 

19.   Amorphous  Sulphur. 

Distil  40  g.  of  sulphur  from  the  bulb  of  a  small  retort  and 
allow  half,  as  it  condenses  in  the  neck,  to  flow  into  a  mixture  of 
ice  and  water,  the  other  half  into  boiling  water.  Dry  these  two 
preparations  externally  with  filter  paper,  then  leave  them  for 
several  hours  at  50°  to  60°  to  dry  still  further.  The  next  day 
extract  5  to  10  g.  of  each  of  the  samples  with  carbon  bisulphide 
in  a  Soxhlet  apparatus,  using  a  weighed  extraction-thimble. 
Determine,  by  weighing  again,  the  percentage  of  each  sample  of 
sulphur  which  has  been  dissolved  in  the  carbon  bisulphide.  The 
sample  which  was  slowly  cooled  dissolves  almost  completely, 
since  it  has  changed  into  the  soluble  rhombic  form;  the  sud- 
denly cooled  sample  leaves  behind  30  to  50%  of  amorphous, 
insoluble  sulphur. 

If  a  mixture  of  sulphur  and  ammonium  carbonate  is  distilled, 
and  the  sulphur,  as  in  the  first  case  above,  is  cooled  suddenly  in  the 
freezing  mixture,  a  product  is  obtained  which  contains  90%  of 
the  soluble  rhombic  form,  since  the  presence  of  ammonia  accel- 
erates catalytically  the  transformation. 

COLLOIDAL  STATE. 

A  large  number  of  substances  are  capable  of  apparently  dissolving  in  water 
to  form  what  may  be  termed  pseudo-solutions;  such  pseudo-solutions  are 
characterized  by  an  extremely  small  diffusive  power,  a  low  osmotic  pressure, 
and  an  inability  to  undergo  dialysis;  in  these  respects  they  differ  from  true 
solutions  of  crystalloids,  and  are  thus  called  colloidal  solutions  (Graham,  1862). 
Colloidal  solutions  are  distinguished,  according  to  the  nature  of  the  solvent, 
as  hydrosols,  alcosols,  etc.  If  the  dissolved,  or,  more  correctly,  the  pseudo- 
dissolved  substance  is  separated  from  the  solvent,  it  is  often  found  not  to 
have  lost  the  power  of  again  passing  into  colloidal  solution  (reversible  colloids). 
On  the  other  hand,  there  are  numerous  substances,  particularly  inorganic  ones, 


34  CHANGES   OF   CONDITION. 

which  are  unable  of  themselves  to  pass  directly  into  colloidal  solution,  but 
can  by  suitable  means  be  brought  into  that  state;  if  such  substances  are 
separated  from  the  solution,  they  do  not  possess  the  power  of  dissolving 
again  unaided  (irreversible  colloidals). 

Solutions  of  irreversible  colloids  can  be  obtained  by  disintegrating  the 
substance  in  the  electric  arc  under  water  (Bredig's  method;  see  No.  20),  by 
forming  the  substance  in  aqueous  solution  by  double  decomposition  and 
preventing  its  precipitation  by  suitable  means  (Nos.  21;  24,  Note  on 
zircon  oxide  hydrosol;  No.  25),  or  by  previously  imparting  to  the  solid  the 
ability  to  dissolve  by  treatment  with  small  quantities  of  alkali  or  acid.  The 
last  method  was  compared  by  Graham  with  the  process  of  digestion  and 
called  by  him  peptonization  (No.  25,  tin-oxide  hydrosol). 

By  means  of  optical  methods  (ultramicroscopic  investigation),  it  has  been 
proved  that  many  of  these  colloidal  solutions  consist  really  of  suspensions  of 
extremely  small  particles  the  size  of  which  can  be  estimated  down  to  a 
diameter  of  about  6  /XAC.  In  this  manner  the  formerly-very-puzzling  fact  of 
the  existence  of  solutions  of  substances  that  are  ordinarily  insoluble,  such  as 
gold,  platinum,  and  the  sulphides  of  the  heavy  metals,  finds  its  explanation. 

Many  irreversible  colloids  separate  out  from  their  solutions  as  volumi- 
nous precipitates  containing  a  large  amount  of  water  (hydrogels}.  Such 
hydrogels  as,  for  example,  those  of  ferric  oxide,  aluminium  oxide,  and  silicon 
oxide,  can  also  be  obtained  directly  by  chemical  precipitation  from  solution. 
These  precipitates  contain  a  great  deal  more  water  than  would  correspond  to 
their  hydroxide  formulas  (Fe(OH),,  A1(OH)8,  Si(OH)4,  etc.),  and  since  the 
water  thus  contained  does  not  show  the  characteristics  of  chemically  com- 
bined water  (compare  the  theoretical  section  on  Hydrates  preceding  No.  144), 
these  precipitates  should  be  designated,  according  to  van  Bemmelen,  as 
oxide  hydrogels,  or  by  their  old  name  of  hydrated  oxides. 

The  change  from  the  hydrosol  to  the  hydrogel  state  is  most  simply 
brought  about  by  the  addition  of  an  electrolyte.  Many  colloidal  solutions 
are  exceedingly  sensitive  to  electrolytes  (gold  solution,  No.  25);  the  most 
essential  condition  for  the  preparation  of  the  solution  must  therefore  be  the 
absence  of  any  unnecessary  electrolyte.  Even  reversible  colloids  can  be 
separated  from  their  solutions  by  the  addition  of  large  amounts  of  an  elec- 
trolyte ;  the  process  is  termed  "  salting  out, "  and  is  employed  in  the  pre- 
cipitation of  proteins  and  dyestuffs  (No.  24). 

ADSORPTION  COMPOUNDS. 

Adsorption  is  the  phenomenon  shown  by  certain  substances  having  a 
large  surface  (for  example,  wood-charcoal)  of  condensing  gaseous  or  dis- 
solved substances  upon  themselves.  Adsorption  is  a  special  kind  of  natural 
phenomenon  which  is  quite  distinct  from  the  process  of  solution. 

It  has  been  shown,  especially  by  van  Bemmelen,  that  the  typical  col- 
loids exhibit  a  particularly  high  power  of  adsorption;  and  also  that  organized 
matter,  such  as  plant  fibers  and  decayed  material  like  humus,  possess  the 
same  power.  The  latter  classes  of  substances  are,  therefore,  included  among 


ADSORPTION  COMPOUNDS.  35 

colloids,  especially  since  in  contrast  to  crystalloids  they  lack  a  definite  form, 
bounded  by  rigid  surfaces. 

Among  other  instances  in  which  adsorption  may  occur,  the  following  are 
of  especial  importance.  It  may  take  place:  (1)  between  liquids  and  solid 
substances  (as  when  moisture  is  retained  by  many  solid  substances  and  par- 
ticularly by  hydrogels);  (2)  between  dissolved  substances  and  solids  (as  in 
the  dragging  down  of  dissolved  salts  by  precipitates  and  in  the  adhering  of 
fertilizer  salts  in  the  soil;  cf.  Nos.  22  and  23);  (3)  between  dissolved  rever- 
sible colloids  and  solid  materials  (as  in  numerous  dyeing  processes;  No.  24). 
The  combinations  so  produced  are  known  as  adsorption  compounds. 

For  many  adsorption  processes,  definite  relations  exist  between  the  con- 
centration in  the  solution  of  the  substance  adsorbed  and  the  composition  of 
the  adsorption  compound.  Adsorption  acts  relatively  more  strongly  the 
less  concentrated  the  solution  (cf.  No.  22).  Many  adsorption  compounds 
are  of  so  specific  a  nature  (for  example,  iodo-starch)  that  it  has  been  only 
after  the  investigation  of  the  quantitative  relations  between  their  com- 
position and  the  conditions  under  which  they  are  formed  that  their 
difference  from  true  chemical  compounds  has  been  established.  Such  an 
exhibition  of  adsorptive  power  as  this  has  also  been  designated  as  affinity  of 
condition. 

Different  dissolved  colloids  can  also  mutually  combine  to  form  adsorption 
compounds:  gold  hydrosol  which  is  so  exceedingly  sensitive  toward  elec- 
trolytes can,  by  the  addition  of  a  non-sensitive  colloid  such  as  gelatin,  be 
itself  made  stable  toward  electrolytes ;  this  would  tend  to  show  that  a  com- 
bination had  taken  place  between  the  two  kinds  of  dissolved  colloids  (cf. 
No.  25).  Substances  which  act  as  gelatin  does  in  this  case  are  known  as 
protective  colloids. 

Again,  other  colloids  are  capable  of  mutually  precipitating  one  another 
out  of  solution.1  Thus  arsenic-sulphide-hydrosol  and  iron-oxide-hydrosol 
when  mixed  in  the  right  proportions  are  both  precipitated  as  a  common 
adsorption  compound  (cf.  No.  24,  4).  This  precipitating  power  arises  from 
the  same  cause  as  another  characteristic  property  which  pseudo-dissolved 
substances  have  in  common  with  suspensions:  if  suspensions  or  colloids  are 
subjected  to  the  action  of  a  strong  electric  potential,  a  passage  of  the  suspended 
material  through  the  solution  occurs,  but  this  is  of  a  very  different  nature 
from  ionic  migration.  (Connective  transference.}  While  with  electrolytes,  the 
dissociated  parts  possess  opposite  electrical  charges,  here  the  opposite  charges 
reside  upon  the  pseudo-dissolved  material  and  the  solvent  itself,  respectively. 
It  is  a  rule  that  two  such  colloids  in  order  to  precipitate  each  other  must 
have  charges  of  opposite  sign  when  referred  to  that  of  the  common  solvent. 
(Example,  zircon-gold-purple,  No.  25.) 

Finally,  adsorption  compounds  can  be  produced  by  precipitating  two 
colloids  out  of  a  common  solution  by  the  addition  of  an  electrolyte.  (Exam- 
ple, purple  of  Cassius,  No.  25.) 


W.  Biltz,  Ber.  37,  1095  (1904) 


36  CHANGES  OF   CONDITION. 

20.    Colloidal  Platinum,  according  to  Bredig. 

Connect  two  platinum  wires  of  1  mm.  diameter,  whose  upper 
ends  are  insulated  by  glass  tubes,  with  the  terminals  of  the  110- 
volt  lighting  circuit,  and  insert  an  ammeter  and  a  resistance. 
Clamp  one  electrode  so  that  its  lower  end  dips  into  100  to  150  c.c. 
of  distilled  water  in  a  glass  dish;  hold  the  other  electrode  in 
the  hand,  and  while  it  is  immersed  in  the  water,  touch  it  to  the 
first  one  and  remove  it  to  such  a  distance  that  a  small  arc  can  be 
maintained.  The  resistance  should  be  regulated  so  that  a  current 
of  6  to  10  amperes  will  flow.1  From  the  cathode  the  disintegrated 
platinum  passes  in  grayish  brown  clouds  into  the  liquid.  It  is  not 
usually  possible  to  preserve  the  arc  for  more  than  a  short  time,  but 
it  can  be  repeatedly  started  again  by  bringing  the  electrodes 
together,  and  then  separating  them;  the  process  may  be  thus  con- 
tinued until  the  platinum  solution  is  so  dark  as  to  be  nearly  opaque, 
or  until  it  has  become  too  hot.  The  platinum  electrodes  should 
never  be  held  except  by  the  insulated  glass  ends.  If  the  experiment 
does  not  succeed  well  at  first,  add  a  trace  of  dilute  alkali  to  the 
water.  Finally,  filter  off  the  coarser  platinum  powder,  and  employ 
the  solution  in  the  following  experiments: 

1.  To  a  few  cubic  centimeters  add  a  drop  of  a  dilute  salt  solu- 
tion.    After  a  short  time  the  colloidal  metal  separates  out  in  the 
form  of  a  powder  (precipitation  by  an  electrolyte). 

2.  Add   some  of   the   platinum  solution  to  dilute    hydrogen 
peroxide.     A  reaction  takes  place  with  evolution  of  oxygen,  while 
a  comparison  sample  of  the  hydrogen  peroxide  itself  remains  clear. 
Concentrated  30  per  cent  hydrogen  peroxide,  when  treated  with 
the  platinum  solution,  decomposes  at  first  slowly,  then  more  rap- 
idly, with  rise  of  temperature,  and  finally  with  explosive  violence. 
This  is  an  instance  of  catalysis;  i.e.,  the  acceleration  by  means  of 
a  chemically  indifferent  substance  (the  platinum)  of  a  reaction 
(the  decomposition  of  hydrogen  peroxide)  which  would  of  itself 
take  place  slowly. 

3.  To  another  portion  of  dilute  hydrogen  peroxide  add  a  few 
drops   of   hydrogen    sulphide   water    and    then   some    platinum 


1  A  technical  ammeter  with  large  capacity  is  used  here.  Before  begin- 
ning the  experiment,  see  that  the  wiring  will  safely  bear  the  current 
required. 


COLLOIDAL  ANTIMONY  SULPHIDE.  37 

solution.  This  time  no  decomposition  occurs;  the  catalyzer  has 
temporarily  lost  its  accelerating  action  in  consequence  of  the 
presence  of  a  third,  likewise  indifferent  substance  which  acts  as; 
a  poison.  The  catalytic  property  of  the  platinum  depends  upon, 
the  condition  of  the  surface  of  the  finely  divided  material;  the 
action  of  the  poison  is  due  perhaps  to  an  alteration  in  the  con- 
dition of  the  surface. 

21.    Colloidal  Antimony  Sulphide. 

Allow  a  cold  1  per  cent  solution  of  tartar  emetic  to  drop  from  a 
dropping  funnel  into  a  solution  of  hydrogen  sulphide  water  through 
which  a  fairly  strong  stream  of  hydrogen  sulphide  is  kept  passing. 
Antimony  sulphide  produced  under  these  conditions  does  not  form 
a  precipitate,  but  remains  in  colloidal  suspension  as  a  deep-orange- 
colored  pseudo-solution,  that  appears  perfectly  clear  by  transmitted 
light.  First  remove  the  excess  of  hydrogen  sulphide  by  a  stream 
of  hydrogen,  then  place  the  solution  in  a  dialyzing  tube  of  parch- 
ment paper,  and  suspend  the  latter  in  a  vessel  of  distilled  water. 
If  the  tube  is  free  from  imperfections,  almost  none  of  the  yellow 
material  passes  into  the  outer  solution,  while  the  salts  present  are 
gradually  removed  by  diffusion.  Replace  the  outside  water  with 
fresh  every  six  hours  at  first,  later  every  twelve  hours,  and  con- 
tinue the  dialysis  about  four  days. 

Dilute  a  part  of  the  antimony  sulphide  solution  to  ten  times  its 
original  volume,  and  carry  out  the  following  experiments : 

1.  Place   roughly   equivalent   normal   solutions   of   potassium 
chloride,  barium  chloride,  and  aluminium  chloride  in  three  burettes 
and  add  each,  drop  by  drop,  to  three  separate  portions  of  20  c,c. 
each  of  the  colloidal  antimony  sulphide  until  the  complete  precipi- 
tation of  the  latter  is  accomplished.     It  may  be  necessary  to  repeat 
the  experiment  before  the  right  end-point  is  obtained,  noticing 
carefully  when  the  precipitant  causes  the  supernatant  solution  to 
change  from  yellow  to  colorless.     It  will  be  found  that  the  largest 
amount  of  precipitant  is  required  in  the  case  of  potassium  chloride, 
the  smallest  in  the  case  of  aluminium  chloride.     The  precipitating 
power  of  equivalent  amounts  of  salts  for  colloids  of  this  class 
increases  with  the  valence  of  the  cation. 

2.  Shake  several  portions  of  10  c.c.  each  of  the  antimony  sul- 
phide for  1  minute  with  2.0,  1.0,  0.5,  0/2,  and  0.1  g.  respectively  of 


38  CHANGES   OF  CONDITION. 

powdered  barium  sulphate.  The  pseudo-dissolved  material  is 
adsorbed  by  the  solid.  Determine  what  quantity  of  the  barium 
sulphate  is  just  sufficient  for  complete  adsorption. 

Test  also,  for  comparison,  the  precipitating  power  of  wood- 
charcoal  and  animal-charcoal. 

22.   Adsorption  of  Iodine  by  Charcoal.     Adsorption  Curve. 

For  a  more  accurate  study  of  adsorption  phenomena,  charcoal 
is  particularly  well  suited,  since  under  ordinary  conditions  it  is 
an  indifferent  substance,  and  has  a  constant  surface.  Place  in 
each  of  four  200  c.c.  glass-stoppered  bottles  2  g.  of  washed,  dried, 
and  ignited  animal  charcoal  and  50,  75,  90,  and  95  c.c.  of  alcohol; 
then  add  50,  25,  10,  and  5  c.c.  respectively  of  a  normal  alcoholic 
iodine  solution,  so  that  the  volume  in  each  bottle  will  be  100  c.c. 
and  the  contents  will  differ  only  in  the  amount  of  iodine.  After 
24  hours,  during  which  time  the  bottles  should  be  frequently 
shaken,  take  10  c.c.  of  the  clear  solution  from  each  and  determine 
the  iodine  content  by  titration  with  0.1-normal  sodium  thio- 
sulphate  solution.  The  difference  between  the  original  weight  pres- 
ent and  the  amount  of  iodine  thus  determined  gives  the  quantity 
adsorbed  by  the  charcoal.  Arrange  the  results  in  a  table  and  plot 
them  on  coordinate  paper  with  the  adsorbed  amounts  as  ordinates 
and  the  corresponding  amounts  left  in  solution  as  abscissas.  It  is 
found  that  from  dilute  solutions  relatively  more  iodine  is  adsorbed 
than  from  concentrated  ones.  This  relation  is  quite  general  for 
adsorption  equilibria,  and  can  frequently  be  represented  by  the 

f  n 

equation  -1-  =  k,  in  which  c1  is  the  concentration  of  the  adsorbed 

C2 

iodine,  c2  that  of  the  unadsorbed  iodine,  and  n  and  k  are  constants. 
Since  in  these  experiments  the  same  volume  and  the  same  amount 
of  charcoal  is  taken  each  time,  the  quantities  of  iodine  determined 
may  be  used  directly  in  place  of  cl  and  c2  in  the  equation.  Find 
by  trial  what  value  of  n  makes  the  quotient  remain  most  nearly 
constant;  n  is  greater  than  1. 

23.   Lanthanum  Blue. 

The  blue  color  which  an  iodine  solution  gives  with  starch  depends  on  the 
formation  of  an  adsorption  compound  between  the  two  substances.  This 
can  be  proved  by  a  quantitative  investigation  of  the  adsorption  curve. 
The  same  conclusion  can  be  reached  from  the  fact  that  another  substance, 
the  hydrogel  of  basic  lanthanum  acetate,  which  has  no  property  in  common 


MOLYBDENUM  BLUE.  39 

with  starch  except  the  colloidal  condition,  gives  precisely  the  same  color 
reaction.  The  formation  of  lanthanum  blue  is  made  use  of  as  a  test  for  lan- 
thanum. (Cf.  No.  171.) 

To  a  dilute  solution  of  lanthanum  acetate,  or  of  lanthanum 
nitrate  acidified  with  acetic  acid,  add  a  solution  of  iodine  in  potas- 
sium iodide,  and  then  introduce  ammonia  cautiously,  so  as  to  not 
quite  cause  the  yellowish-brown  color  of  the  iodine  to  disappear; 
warm  very  gently,  and  a  dark-blue  precipitate  gradually  forms. 
If  a  very  dilute  solution  of  lanthanum  acetate  is  used,  a  blue 
colloidal  solution  is  produced  instead. 

Prepare  also  a  little  iodo-starch  from  a  very  dilute  starch  solu- 
tion and  a  few  drops  of  iodine  solution,  and  compare  the  colors. 

24.   Molybdenum  Blue,  Mo3O8. 

Blue  molybdenum  oxide,  Mo3O8,  is  soluble  as  a  reversible  colloid  in  water 
and  behaves  in  this  respect,  as  well  as  in  its  behavior  towards  vegetable  and 
animal  fibers,  like  many  of  the  organic  dyestuffs. 

Dissolve  15  g.  of  commercial  ammonium  molybdate, 
5  (NH4)2MoO4  •  7  MoO3  •  7  H2O,  in  250  c.c.  of  water  and  35  to  40  c.c. 
of  2-normal  sulphuric  acid,  heat  the  solution  to  boiling,  and  keep 
boiling  gently  for  30  minutes  to  an  hour,  meanwhile  introducing  a 
stream  of  hydrogen  sulphide.  Reduction  takes  place,  causing  the 
appearance  in  a  few  moments  of  a  dark-blue  color.  With  a  smaller 
amount  of  sulphuric  acid  a  precipitation  of  molybdenum  sulphide 
takes  place;  with  more  acid  the  yield  becomes  poorer.  Filter  from 
precipitated  sulphur  and  subject  the  solution  to  dialysis  for  from 
4  to  6  days  (cf.  No.  21),  until  the  outside  water  is  free  from  sul- 
phuric acid,  and  is  only  faintly  blue.  Evaporate  the  contents  of 
the  dialyzing  tube  in  a  porcelain  dish,  at  first  over  a  free  flame  and 
finally  on  the  water  bath,  until,  after  frequent  stirring,  the  resinous, 
deep-blue  residue  has  become  a  dry  powder.  Yield,  5  to  7  g. 

The  preparation  is  soluble  without  residue  in  water.  Use  the 
solution  for  the  following  experiments: 

1.  Boil  a  portion  with  a  piece  of  undyed  silk.     The  silk  is  col- 
ored blue.     This  dyeing  experiment  can  also  be  carried  out  with 
the  undialyzed  solution. 

2.  Repeat  Experiment  1  with  the  addition  of  a  considerable 
amount  of  sodium  sulphate.     More  of  the  coloring  matter  is  taken 
up  by  the  silk.     This  is  due  to  the  salting-out  action  of  the  elec- 
trolyte. 


40  CHANGES    OF    CONDITION. 

3.  Shake  a  little  of  the  solution  with  some  freshly-precipitated, 
washed  aluminium  hydroxide.     A  mixture  of  molybdenum  blue 
and  hydrated  aluminium  oxide  is  precipitated  (a  lake)  and  the 
solution  becomes  lighter  colored.     Hydrogels  can  thus  be  dyed 
like  fabrics. 

4.  Mix  a  few  cubic  centimeters  of  dilute  molybdenum  blue  solu- 
tion, drop  by  drop,  with  a  colloidal  solution  of  zircon  oxide.1    After 
enough  of  the  former  has  been  added,  a  molybdenum-blue-zircon 
lake  is  precipitated.     Regarding  the  mutual  precipitation  of  col- 
loids, see  p.  35. 

25.    Colloidal  Gold  Solutions;  Precipitating  Colloids  and  Protective 

Colloids. 

During  recent  years,  gold  solutions  have  been  used  with  great  success, 
particularly  by  Zsigmondy,  in  making  clear  the  nature  of  the  various  prop- 
erties of  colloidal  suspensions.  Red  colloidal  solutions  of  gold  can  be  pre- 
pared by  electrical  disintegration  of  the  metal  or  by  reduction  of  gold  salts 
in  various  ways.  The  red  color  of  gold-ruby  glass  is  likewise  due  to  colloidal 
gold.  The  adsorption  compound  of  colloidal  gold  and  tin  oxide  hydrogel, 
called  purple  of  Cassius,  serves  for  the  qualitative  detection  of  gold. 
(Cf.  No.  9.) 

1.  Colloidal  Gold  according  to  Zsigmondy.2  Redistil  120  c.c.  of 
ordinary  distilled  water,  using  a  silver  condenser,3  and  a  receiver  of 
Jena  glass.  Heat  the  distillate  to  boiling  in  a  300  to  400  c.c.  Jena 
flask,  adding  during  the  heating  2.5  c.c.  of  a  solution  of  chlorauric 
acid  (0.6  g.  of  crystallized  HAuCl4.3  H2O  in  100  c.c.  of  water)  and 
3.0  to  3.5  c.c.  of  an  0.18-molal  solution  of  purest  potassium  car- 
bonate. Immediately  after  the  solution  boils,  add,  while  vigor- 
ously rotating,  3  to  5  c.c.  of  a  dilute,  freshly-distilled  solution  of 
formaldehyde  (0.3  c.c.  of  the  commercial,  distilled  formalin  in 
100  c.c.  of  water).  This  solution  should  be  added  from  a  pipette 
in  several  portions,  best  with  the  flame  removed,  waiting  each 
time  a  few  seconds,  or  at  the  most  a  minute,  for  the  reaction  to 
take  place.  A  light-red  color  appears,  which  changes  to  an  intense 
bright  red  in  a  few  seconds;  the  latter  color  persists  unchanged. 


1  Pseudo-solutions  of  zircon  oxide  can  be  prepared  by  dialyzing  a  solu- 
tion (about  16  per  cent)  of  zircon  nitrate  for  about   5  days  with  frequent 
change  of  water. 

2  R.  Zsigmondy,  Z.  Anal.  Chem.  40,  697  (1901).    The  most  beautiful  gold 
solutions  can  be  obtained  by  this  classic  but  very  difficult  procedure. 

3  The  use  of  a  silver  condenser  is  essential. 


COLLOIDAL  GOLD  SOLUTIONS.  41 

If  violet  or  blue  'solutions  appear  at  first,  the  proportions  of 
formaldehyde  and  carbonate,  and  the  method  of  heating  should  be 
altered.  This  method  of  preparation  is  exceedingly  sensitive  to 
slight  changes  in  the  conditions  of  the  procedure;  it  does  not  yield 
uniformly  good  results  in  all  laboratories. 

This  colloidal  gold  solution  can  be  obtained  with  more  certainty 
even  with  ordinary  distilled  water,  if  in  accordance  with  a  new 
method  of  Zsigmondy,  the  solution  is  inoculated  before  its  reduc- 
tion with  a  little  of  another  deep-red  colloidal  gold  solution  which 
has  been  reduced  by  means  of  phosphorus.  This  solution  can  be 
prepared  exactly  as  directed  above,  except  that  for  the  reducing 
agent  a  few  drops  of  a  five  times  diluted  saturated  solution  of  dry 
phosphorus  in  ether  is  used.  This  should  be  added  before  warm- 
ing, then,  on  heating,  the  solution  becomes  colored  a  deep  red. 
Boiling  should  be  continued  until  the  odor  of  ether  has  disappeared. 
The  path  of  a  beam  of  light  passing  through  the  solution  should  be 
either  not  at  all  or  scarcely  visible  when  observed  from  the  side, 
although  on  the  other  hand  the  so-called  "  Tyndall  phenomenon  " 
is  very  characteristic  for  the  great  majority  of  colloidal  solutions. 

If  a  few  drops  of  the  solution  prepared  as  just  described  are  added 
as  an  inoculating  fluid  to  the  gold  solution  before  its  reduction  with 
formaldehyde  by  the  above  process,  good  results  are  then  quite 
certain  to  be  obtained. 

2.  Colloidal  Gold  according  to  Donau.     Prepare  carbon  monox- 
ide by  heating  together  oxalic  and  concentrated  sulphuric  acids 
and   passing  the  evolved   gases    through   sodium   hydroxide   to 
remove  the  carbon  dioxide;  conduct  a  slow  current  of  the  gas 
through  120  c.c.  of  a  0.03  per  cent  chlorauric  acid  solution.    There 
is  produced  first  a  violet,  then  a  violet-red,  and  later  a  deep-red 
color;  do  not  carry  the  reduction  beyond  this  point,  or  a  violet 
coloration  will  be  produced. 

3.  Colloidal  Gold  according  to  Brunch.    Reduce  a  boiling  solu- 
tion of  chlorauric  acid,  of  the  same  quantity  and  concentration 
as  given  in  the  last  paragraph,  by  adding  1  to  3  c.c.  of  a  0.2  per 
cent  solution  of  sodium  hyposulphite.1 

Use  a  gold  solution  prepared  by  any  one  of  these  methods  in  the 
following  experiments. 

1  This  refers  to  true  sodium  hyposulphite,  Na,>S2O4,  not  sodium  thiosulphate 
which  is  often  called  "  hypo." 


42  CHANGES   OF   CONDITION. 

I.  Synthesis  of  Purple  of  Cassius.     From  a  mixture  of  colloidal 
stannic  oxide  and  colloidal  gold,  the  gold-stannic  oxide  adsorption 
compound  is  precipitated  on  the  addition  of  electrolytes.     In  order 
to  prepare  the  colloidal  stannic  acid  solution,  allow  5  c.c.  of  tin 
tetrachloride  (No.  50)  to  become  hydrolyzed  by  the  addition  of 
150  to  200  c.c.  of  water,  and  pour  this  solution  into  500  c.c.  of 
water  to  which  a  few  drops  of  ammonia  have  been  added.  Dialyze 
the  clear  mixture  for  five  days  (cf.  No.  21),  changing  the  outside 
water  two  or  three  times  daily,  until  it  shows  no  test  for  chlorides. 
If  during  this  process  a  hydrogel  separates  in  the  dialyzing  tube, 
it  may  be  peptonized  (see  p.  34)  in  a  beaker  by  the  addition  of 
about  three  drops  of  ammonia,  whereupon,  after  a  time,  the  jelly 
will  go  over  into  a  perfectly  clear  hydrosol.     A  mixture  of  this 
hydrosol  with  an  equal  volume  of  gold  solution  remains  unchanged, 
but  on  addition  of  a  salt  (ammonium  chloride)  a  beautiful  deep- 
reddish-purple  precipitate  is  formed  which  can  be  filtered  off;  the 
compound  is  characterized  by  its  solubility  in  ammonia. 

II.  Zircon-gold-purple.     Treat  130  c.c.  of  a  boiling  colloidal  gold 
solution  with  28  c.c.  of  a  boiling  colloidal  zircon  solution  (see  foot- 
note to  No.  24).     A  precipitation  of  zircon-gold-purple  takes  place 
even  without  the  addition  of  an  electrolyte.     In  the  cold  the  pre- 
cipitate forms  slowly.     If  the  deposition  is  incomplete,  preliminary 
tests  must  be  made  with  small  portions  to  find  the  right  propor- 
tions in  which  to  mix  the  above  solution. 

III.  Gold  Solution  and  Protective  Colloids.   Treat  10  c.c.  of  col- 
loidal gold  solution  with  a  few  drops  of  dilute  hydrochloric  acid;  a 
blue  coloration  is  first  produced,  later  sedimentation  of  the  metal. 

Repeat  the  experiment  after  first  adding  one  drop  of  a  dilute 
gelatin  solution  (0.2  per  cent)  to  the  colloidal  gold;  no  change 
whatever  in  the  color  or  the  stability  of  the  gold  solution  is  observed. 

26.   Hydrogels  as  Semipermeable  Membranes. 

The  separation  of  colloids  and  electrolytes  by  dialysis  depends  on  the 
colloidal  nature  of  the  parchment  wall,  which  is  impervious  to  other  colloids. 
•Certain  colloids  are  impervious  even  to  truly  dissolved  substances  but  still 
pervious  to  water.  By  means  of  membranes  of  such  semipermeable  material 
a  solute  can  be  separated  from  its  solvent,  and  thus  the  osmotic  pressure  of 
the  dissolved  substance  can  be  both  demonstrated  and  measured.  Cupric 
ferrocyanide  has  been  found  especially  suitable  as  a  semipermeable  colloid. 
(Pfeffer,  1877.) 


SEMIPERMEABLE   MEMBRANES.  43 

1.  Cupric  Ferrocyanide  Membrane.     Let  a  drop  of  a  cold,  satu- 
rated  potassium   ferrocyanide   solution   run   from   a   fine,    glass 
capillary  into  a  0.5-molal  copper  sulphate  solution,  and  detach  it 
by  means  of  a  slight  motion,  so  that  it  sinks  to  the  bottom  of  the 
vessel.     The  drop  has  at  the  moment  of  its  entrance  into  the  solu- 
tion become  surrounded  with  a  thin  film  of  cupric  ferrocyanide, 
which  keeps  growing  at  the  cost  of  the  dissolved  components. 
Since,  however,  the  concentration  of  the  solute  within  the  mem- 
brane is  greater  than  that  of  the  copper  sulphate  outside,  the  mem- 
brane expands  in  consequence  of  the  pressure  caused  by  the  water 
entering  through  the   walls.     The   membrane   is   at   first   trans- 
parent and  traversed  by  brown  veins.     A  uniform  growth  can  be 
brought  about  by  occasional,  gentle  stirring  of  the  copper  sulphate 
solution.     As  the  cell  keeps  expanding  a  point  is  reached  where, 
the  specific  gravity  of  its  contents   having   grown  less  through 
entrance  of  water,  the  cell  rises  to  the  surface  of  the  solution,  and 
remains  there  until,  after  ten  or  fifteen  minutes,  the  constant 
thickening  of  its  walls  so  increases  its  weight  as  to  make  it  once 
more  sink,  and  this  time  permanently. 

2.  Membranes  of  Colloidal  Silicates  of  the  Heavy  Metals.     Dilute 
some  commercial  water-glass  solution  until  a  specific  gravity  of  1.1 
is  obtained.     Into  about  100  c.c.  of  this  solution  in  a  narrow  beaker, 
drop  small  particles  of  various  salts,  such  as  copper  sulphate, 
aluminum  sulphate,  ferric  chloride,  nickel  nitrate,  cobalt  nitrate, 
manganous  sulphate,  lead  nitrate,  and  uranyl  nitrate.     Within  a 
few  minutes  the  particles  begin  to  swell  and  to  send  out  shoots 
which  branch  and  grow  toward  the  surface  of  the  liquid  until  the 
whole  beaker  is  filled  with  what  appears  like  bright  colored  alga? 
growths.     The  salt,  on  being  thrown  into  the  solution,  begins  at 
once  to  dissolve,  and  at  the  surface  of  contact  between  this  solu- 
tion and  the  silicate  solution,  an  insoluble  semi-permeable  film  of 
metal  silicate  is  formed.     The  dissolved  salt  within  exerts  an 
osmotic  pressure  against  this  film,  and  forces  it  to  expand,  while 
the  water  which  is  thereby  drawn  in  through  the  film  dissolves 
more  of  the  salt.     The  osmotic  pressure  is  thus  maintained  and 
the  film  is  continuously  forced  to  expand  until  it  bursts  in  places 
and  forms  outgrowths  and  side-arms. 


CHAPTER  III. 


SIMPLE   COMPOUNDS. 

UNDER  the  designation  simple  compounds  are  included  all  compounds  con- 
taining but  two  elements,  with  the  exception,  however,  of  the  persulphides, 
peroxides,  polyhalides,  etc.,  which  are  considered  to  have  complex  cations 
(Chap.  IV).  The  metal  hydroxides  and  cyanides,  in  which  the  radicals  OH 
and  CN  behave  as  single  elements,  are  also  classed  as  simple  compounds. 

METHODS  OF  PREPARATION.     Simple  compounds  are  prepared: 

(1)    Synthetically  from  the  elements: 

Br2  +  H2  =  2HBr(No.  35a). 

Cf .  also  Cerium  Hydride  (No.  32) ;  FeI2  (No.  39) ;  FeCl3  (No.  42) ;  CrCl3  (No.  44) ; 
S2C12  (No.  45) ;  PC13  (No.  46) ;  BiI3  (No.  48) ;  BiBr3  (No.  49) ;  SnCl4  (No.  50) ; 
SiCl4  (No.  51);  P2S5  (No.  54);  HgS  (No.  55) ;  Mg3N2  (No.  61);  Mg3P2  (No.  63). 

Frequently  the  synthetic  preparation  of  compounds  between  two  elements 
takes  place  in  stages: 

P  +  3  Cl  =  PC13, 
PC13  +  2  Cl  =  PC15.    (No.  46) 

Cf.  also  SO3  (No.  28);  SbCl5  (No.  47);  SnS2  (No.  56). 

(2)  By  the  interaction  of  two  substances,  each  containing  one  of  the  ele- 
ments which  are  to  be  combined : 

(a)    By  the  action  of  an  element  upon  a  compound: 

Fe  +  2  HC1  =  FeCl2  +  H,.     (No.  43) 
Cf.  also  A1C13  (No.  43). 

(6)    By  the  chemical  reaction  between  two  compounds  (double  decomposi- 
tion or  metathesis),  as,  for  example,  in  the  precipitation  of  a  sulphide  from  the 
solution  of  a  metallic  salt  by  means  of  hydrogen  sulphide : 
CuCl2  +  H2S  =  CuS  +  2  HC1. 

Cf.  MnS  (No.  57);  TiS2  (No.  58);  CrN  (No.  62);  and  BrH  (No.  35b). 

(3)  By  the  breaking  down  of  more  complicated  compounds: 

2  HNO3  =  H2O  +  N2O4  +  O  (No.    29) 

Cf.  also  Cr2O3  (No.  30);  Cu2O  (No.  31);  copper  hydride  (No.  33);  cyanogen 
(No.  59);  BaS  (No.  87). 

Many  simple  compounds  which  were  formerly  prepared  by  double  decom- 
position, or  by  the  breaking  down  of  more  complicated  compounds,  are  now 
most  advantageously  obtained  directly  from  the  elements,  even  industrially; 
both  because  the  requisite  conditions  for  their  formation  are  now  better 
understood  and  because  some  elements  are  far  more  accessible  than  formerly. 

44 


SIMPLE  COMPOUNDS.  45 

Thus  sulphuric  acid  anhydride  is  now  prepared  directly  by  the  contact  process 
(No.  28),  nitric  oxide  from  the  elements  in  the  atmosphere  (combustion  of  air), 
and  aluminium  chloride  from  the  metal. 

REACTIVITY  AND  DEGREE  OF  DISSOCIATION.  It  is  a  general  principle  that 
the  reactivity  of  a  substance  is  determined  by  a  previous  breaking  down 
(dissociation),  to  a  greater  or  less  extent,  in  the  same  sense  as  that  in  which 
the  reaction  in  question  takes  place.  Phosphorus  pentachloride,  for  example, 
has  a  chlorinating  effect,  and  sulphur  trioxide  an  oxidizing  effect,  only  when 
under  the  prevailing  conditions  the  one  is  partly  dissociated  into  free  chlorine, 
the  other  into  free  oxygen,  even  in  the  absence  of  any  substance  to  be  chlo- 
rinated or  oxidized.  Conversely,  a  tendency  shown  by  substances  to  enter 
into  reaction  may  be  considered  as  an  indication  of  the  preexistence  of  a 
corresponding  dissociation.  The  dissociation  of  binary  substances,  which  in 
fact  often  takes  place  in  stages,  is  essentially  the  reverse  of  their  synthesis. 

PC15  =  PC13  +  C12 
PC13  =  P  +  3  Cl 
2  SO3  =  2  SO2  +  O2     (No.  28  ) 

DISSOCIATION  AND  STATE  OF  EQUILIBRIUM.  If  it  is  true  that  reactivity  is 
dependent  upon  a  certain  ability  to  dissociate,  it  becomes  important  to  study 
the  conditions  favoring  the  formation  and  those  favoring  the  decomposition 
of  substances.  To  take  a  concrete  example,  —  When  does  the  reaction 

2SO2  +  O2  -  2SO3 
take  place,  and  when 

2  SO3  =  2  SO2  +  O2? 

In  this  connection,  another  principle  which  is  likewise  of  very  general 
importance  has  been  established,  —  namely,  that  a  reaction  never  takes  place 
completely  in  one  direction;  at  most  the  chemical  change  may  proceed  chiefly 
in  a  definite  direction  until  when  the  reaction  comes  to  a  standstill  (i.e.,  when 
equilibrium  is  reached)  the  products  of  dissociation  and  the  undissociated 
compound  exist  together  side  by  side,  forming  the  so-called  equilibrium- 
mixture.  The  percentage  composition  of  an  equilibrium-mixture  is  charac- 
terized by  the  fact  that  the  same  values  are  obtained  irrespective  of  whether 
at  the  start  a  mixture  of  the  pure  components  or  the  pure  compound  itself  is 
present.  Thus,  for  example,  the  same  mixture  of  SO3,  SO2,  and  O2  is  obtained 
whether  equivalent  amounts  (e.g.,  formula  weights)  of  SO2  and  O2  are  allowed 
to  react,  or  an  equivalent  quantity  of  SO3  is  allowed  to  decompose  under  the 
same  conditions  of  temperature  and  volume. 

2  SO3  <=±  2  SO2  +  O2 

should  be  read:  sulphur  trioxide  "in  equilibrium  with1"  sulphur  dioxide  and 
oxygen. 

Inasmuch  as  all  reactions  are,  strictly  speaking,  reversible,  it  is  theoretically 
impossible  to  prepare  perfectly  pure  compounds;  for  a  compound  can  only 
exist  as  a  stable  substance  when  it  is  in  equilibrium  with  its  products  of  disso- 
ciation. For  practical  purposes,  however,  it  is  true  that  (1)  in  the  equilibrium 
mixture  the  percentage  content  of  dissociation  products,  or  in  the  other  case 
the  fraction  of  undissociated  substance,  is  frequently  so  extremely  small  that  it 


46  SIMPLE  COMPOUNDS. 

becomes  negligible,  and  that  (2)  the  rate  at  which  certain  substances  decom- 
pose after  they  have  once  been  prepared  "pure"  is  often  so  extremely  slow 
that  measurable  quantities  of  the  dissociation  products  are  formed  only  after  a 
very  long  time.1 

MASS-  ACTION  LAW.  The  state  of  equilibrium  which  combining  or  decom- 
posing substances  reach  is  dependent  (1)  upon  the  nature  of  the  reacting 
substances,  (2)  upon  their,  masses,  and  (3)  upon  the  temperature.  The  influ- 
ence exerted  by  the  masses  upon  the  state  of  equilibrium  can  be  expressed 
mathematically  by  the  so-called  Law  of  Mass  Action  (Guldberg  and  Waage, 
1867).  According  to  this  law,  the  product  of  the  concentrations  of  the  sub- 
stances which  are  upon  the  right-hand  side  of  the  sign  of  equilibrium,  divided 
by  the  product  of  the  concentrations  of  the  substances  on  the  left-hand  side, 
is  a  constant  at  a  given  temperature.  The  concentration  is  usually  expressed 
as  the  number  of  gram-molecules  of  substance  which  are  contained  in  a  unit 
of  volume.  If  A  is  the  formula  of  a  substance,  it  is  customary  to  express  the 
concentration  of  A  by  inclosing  it  in  brackets  [A];  then  if  (A)  represents  the 
actual  amount  of  substance  present  expressed  in  gram-molecules,  and  v 
the  volume,  we  have  /  A  \ 

[A]  -<£ 

If  A  and  B  are  two  substances  which  by  reacting  together  form  two  new- 
substances  C  and  D,  with  which  they  finally  come  to  equilibrium, 

A  +  B^C  +  D, 
then  the  mass-action  law  is  expressed  as  follows: 

[C]  [D]  _  „ 

[A]  [B] 

If,  however,  two  or  more  molecules  (a,  6,  etc.,  being  the  numbers)  of  any  of 
the  substances  enter  into  the  reaction,  then  the  concentration  of  these  sub- 
stances must  be  taken  a,  6,  etc.,  times  in  the  mass-action-law  equation.  Thus 
if  a  molecules  of  A  react  with  b  molecules  of  B  to  form  c  molecules  of  C  and  d 
molecules  of  D,  the  equation  becomes: 


[A]a  [B]6 

APPLICATION  OF  THE  MASS-ACTION  LAW.  The  value  of  the  mass-action 
law  for  the  manufacturing  chemist  becomes  apparent  when  with  its  aid  the 
yields  are  predicted  that  can  be  obtained  in  the  preparation  of  a  substance  at 
a  given  temperature  but  with  varying  proportions  of  the  reacting  materials. 
This  is  particularly  well  illustrated  by  measurements  of  Bodenstein  and  Pohl 
with  regard  to  the  contact-process  for  the  manufacture  of  sulphuric  acid. 
Sulphur  dioxide  and  oxygen  react  within  a  reasonable  interval  of  time  to  form 
sulphur  trioxide  only  when  in  the  presence  of  catalyzers;  the  presence  of  the 
catalyzer,  however,  has  no  effect  upon  the  equilibrium  which  is  finally  reached. 
_  2  SO2  +  O2  +±  2  SO3. 

1  In  old  collections  of  organic  preparations,  the  amount  of  impurities  which 
have  arisen  from  a  self-decomposition  of  the  material  is  often  very  considerable. 


SIMPLE    COMPOUNDS. 


47 


If,  as  before,  v  represents  the  volume  of  the  gas-mixture,  and  the  formulas 
inclosed  in  parentheses  the  number  of  gram-molecules  of  substance  present 
when  equilibrium  is  reached,  then  for  a  given  temperature  the  mass-action 
law  gives  the  relation: 


(S02)2  (02) 
v-        v 


K, 


and  the  "efficiency"  of  the  reaction,  i.e.,  the  ratio  of  the  trioxide  formed  to  the 

unchanged  dioxide,  is:  

(SO3)  _     /K  (Oo) 

(soi)~  V  -y.? 

It  is  evident  from  this  last  expression,  that  it  is  favorable  to  the  yield  if  as 
pure  oxygen  as  possible  (small  v,  little  diluent  of  indifferent  gas)  and  as  much 
oxygen  as  possible  (high  concentration  of  O2)  is  present.  That  this  conclusion 
is  correct  is  shown  by  the  following  table,  in  which  the  yield  (i.e.,  the  actual 
quantity  of  SO3  obtained,  compared  with  what  would  result  if  the  entire 
amount  of  the  SO2  could  be  oxidized)  is  given  in  per  cent  by  volume.  In  the 
first  case  a  mixture  composed  of  the  theoretically  correct  proportions  of  sul- 
phur dioxide  and  oxygen  was  taken,  in  the  second  the  same  mixture  diluted 
with  nitrogen,  and  in  the  third  case  sulphur  dioxide  together  with  an  excess  of 
oxygen.  The  measurements  were  made  at  500°. 


Per  cent  N2. 

Per  cent  SO2. 

Per  cent  O2. 

Yield  in  per  cent 
SO2  oxidized. 

1 

0. 

66.67 

33.33 

91.3 

2 

89.50 

7.00 

3.5 

81.2 

3 

0. 

7.00 

93.00 

98.1 

Further  applications  of  the  mass-action  law  are  illustrated  in  the  prepara- 
tion of  nitrogen  peroxide  (No.  29),  hydrobromic  and  hydriodic  acids  from  the 
elements  (No.  35),  and  phosphorus  pentachloride  and  trichloride  (No.  46); 
cf.  also,  Dissociation  of  Electrolytes,  p.  59. 

DEPENDENCE  OF  EQUILIBRIUM  CONSTANTS  UPON  THE  TEMPERATURE.  It  is 
also  apparent  from  the  above  expression  that  the  yield  in  the  contact  process 
is  dependent  on  the  value  of  the  constant,  K;  if  K  can  be  made  greater,  the 
proportion  of  SO3  is  increased.  The  value  of  K  depends  upon  the  temperature, 
and  its  variation  can  be  predicted  with  the  aid  of  thermochemical  data  and 
the  principles  of  thermodynamics. 

Chemical  reactions  are,  from  a  thermochemical  standpoint,  divided  into 
two  classes:  those  in  which  heat  is  evolved,  or  set  free  (exothermic  reactions), 
and  those  in  which  heat  is  absorbed,  or  used  up  (endothermic  reactions). 
Endothermic  compounds,  or,  in  other  words,  those  in  the  formation  of  which 
from  their  elements  heat  is  absorbed,  are  far  less  common  (cf.  cyanogen, 


48 


SIMPLE    COMPOUNDS. 


No.  59,  and  hydrogen  peroxide,  No.  67).  Experience  has  shown  that  the 
conditions  for  the  formation  of  endothermic  compounds  are  more  favorable 
at  high  temperatures,  while  for  exothermic  compounds  the  reverse  is  true. 
Since  in  the  formation  of  a  substance  the  size  of  the  constant,  K,  is,  under 
otherwise  equal  conditions,  a  measure  of  the  yield,  it  seems  plausible  that  in 
the  case  of  exothermic  reactions  the  value  of  K  diminishes  with  rise  of  tem- 
perature, whereas  in  endothermic  reactions  it  increases. 

According  to  van't  Hoff,  there  exists  between  the  heat  of  reaction  Q,  the 
absolute  temperatures,  Tl  and  T2,  at  which  the  reaction  takes  place,  and  the 
corresponding  equilibrium-constants  K^  and  K2,  together  with  the  gas  con- 
stant R,  the  exact  relation : 

Q 


lnK2  -  InK, 


Q  fl  _  11 
R  LT2      Tj 


from  which  it  follows  that  when  T2  >  Tv  and  the  value  of  Q  is  positive  (exo- 
thermic reactions),  K2  becomes  smaller  than  Kv  and  for  negative  values  of  Q 
(endothermic  reactions)  K2  is  greater  than  K^  This  is  based  upon  the  assump- 
tion that  Q  is  independent  of  the  temperature  at  which  the  reaction  takes 
place.  Whether  this  is  true  or  not  must  be  ascertained  in  the  case  of  each 
reaction  studied.  In  the  synthesis  of  sulphuric  anhydride  this  has  been  found 
to  be  practically  true;  and  it  is  therefore  possible,  when  the  analysis  of  the 
equilibrium  mixture  at  a  given  temperature  is  known,  to  compute  the  com- 
position of  the  equilibrium  mixture  at  any  other  temperature.  The  equi- 
librium-constants, as  above  defined,  for  the  sulphuric  acid  contact-process 
have  been  found  to  be  as  follows: 


/. 

K. 

t. 

K. 

528° 
579 
627 
680 

645.       X102 
131.       X102 
31.6     X102 
8.93  X102 

727 

789 
832 
897 

2.82  X102 
0.794X102 
0.357X102 
0.123X102 

In  accordance  with  the  positive  heat  of  reaction,  Q  =  21,700  calories,  a 
rapid  diminution  in  the  value  of  K  is  observed. 

The  practical  application  of  the  theory  in  the  manufacture  of  sulphuric  acid 
is  shown  by  the  following  yields  calculated  for  various  mixtures  at  different 
temperatures: 


Composition  of  Reacting  Mixtures. 


Yield  of  SO3  at 


S02 

°2 

400° 

500° 

700° 

900° 

66.67 

33.33 

98.1 

91.3 

51.5 

16.0 

14.00 

86.00 

99.8 

97.9 

69.8 

24.4 

2.00 

98.00 

99.8 

98.2 

71.2 

25.6 

It  is,  therefore,  more  important  that  the  process  should  be  carried  out  at  a 
relatively  low  temperature  than  that  an  excess  of  oxygen  should  be  employed. 


LIQUID  SULPHUR  DIOXIDE.  49 

On  the  other  hand,  the  temperature  cannot  be  made  too  low,  as  then  the  rate 
at  which  the  reaction  takes  place,  even  in  the  presence  of  a  catalyzer,  becomes 
too  small. 

The  so-called  "blast-furnace  equilibrium,"  2  CO  +  O2  <=±  2  CO2,  is  displaced 
with  increase  of  temperature,  and  the  reaction  proceeds  more  in  the  direction 
from  right  to  left,  because  here  again  the  heat  of  reaction  is  positive;  cf.  No.  1. 

OXIDES. 
27.    Liquid  Sulphur  Dioxide:    Critical  Point. 

Sulphur  dioxide  is  prepared   technically  by  burning  either  sulphur  or 
pyrite.     On  a  small  scale,  it  is  obtained  by  the  reduction  of  concentrated 
sulphuric  acid,  or  its  anhydride  which  when  hot  has  a  strong  oxidizing  power: 
SO3  +  Cu  =  SO2  +  CuO. 

The  reduction  with  copper  takes  place  also  to  a  slight  extent,  according  to  the 
equation : 

SO3  +  4  Cu  =  CuS  +  3  CuO. 

The  copper  oxide,  as  fast  as  it  is  formed,  dissolves  in  the  sulphuric  acid, 
forming  copper  sulphate.  It  is  essential  for  the  decomposition  that  the 
sulphuric  acid  should  be  hot  and  that  it  should  be  concentrated,  whereby  its 
content  of  SO3  is  increased;  the  nature  of  the  reducing  agent  is  less  important, 
for  the  copper  may  be  replaced  by  other  metals  or  even  by  carbon. 

Sulphur  dioxide  can  be  condensed  to  a  liquid  (boiling-point  —  10°)  by 
cooling  the  gas  in  a  mixture  of  ice  and  common  salt. 

Heat  50  g.  of  copper  turnings  in  a  round-bottomed  flask  with 
200  g.  of  concentrated,  commercial  sulphuric  acid  until  the 
boiling-point  of  the  latter  is  nearly  reached.  Lower  the  flame  as 
soon  as  gas  is  given  off  freely.  Pass  the  gas  through  a  wash 
bottle  containing  concentrated  sulphuric  acid,  and  into  a  second 
empty  wash  bottle  which  is  surrounded  with  a  mixture  of  three 
parts  of  ice  to  one  of  salt,  and  in  which  the  sulphur  dioxide  con- 
denses to  a  liquid. 

When  the  evolution  of  the  sulphur  dioxide  slackens,  pour  the 
liquid  from  the  evolution  flask  into  an  evaporating  dish  before 
it  has  a  chance  to  solidify,  and  allow  it  to  cool  by  standing  over 
night.  In  the  morning  decant  the  liquor  from  the  mass  of 
crystals  which  have  separated,  dissolve  the  crystals  in  as  little 
boiling  water  as  possible,  and  filter  off  any  insoluble  black 
powder  on  a  large  plaited  filter.  Copper  vitriol,  CuS04«  5  H2O 
separates  from  the  filtrate  in  well-defined  crystals;  collect  them 
in  a  filter  funnel  and  evaporate  the  mother  liquor  to  obtain  a 
second,  and  finally  a  third,  crop  of  crystals.  When  the  product 


50 


OXIDES. 


has  been  dried  as  much  as  possible  by  suction,  place  it  in  an 
evaporating  dish,  which  is  covered  with  filter  paper,  and  allow  it 
to  dry  for  several  days  at  the  room  temperature. 

The  sulphur  dioxide,  as  prepared  above,  always  contains  some 
sulphuric  acid,  fumes  of  which  are  carried  over  mechanically 
by  the  gas  from  the  evolution  flask,  and  are  not  entirely  kept 
back  by  the  first  wash-bottle;  immediately  after  being  pre- 
pared, therefore,  the  sulphur  dioxide 
should  be  purified  by  distillation.  Close 
one  tube  of  the  wash  bottle  which 
contains  it,  and  connect  the  other  tube 
by  means  of  a  short  piece  of  rubber 
with  a  glass  tube  which  is  bent  at  right 
angles,  and  whose  vertical  arm  is  drawn 
out  to  an  internal  diameter  of  about 
0.2  cm.  (Fig.  12).  Introduce  this  nar- 
row tube  nearly  to  the  bottom  of  a 
thick-walled  sealing  tube,  0.4  to  0.5  cm. 
wide,  which  is  sealed  at  the  bottom, 
and  drawn  out  a  little  at  a  point  about 
18  cm.  above  the  lower  end.  This 
tube  is  prepared,  cleaned,  and  dried 
before  beginning  the  distillation;  care 

must  be  taken  to  round  the  lower  end  and  to  make  the  con- 
striction without  lessening  the  thickness  of  the  walls  at  any 
point. 

In  order  to  distil  the  sulphur  dioxide,  transfer  the  wash  bottle 
from  the  freezing  mixture  to  a  bath  of  water  at  room  tempera- 
ture, and  place  the  thick-walled  tube,  which  now  serves  as  the 
receiver,  in  the  freezing  mixture.  The  sulphur  dioxide  soon 
begins  to  distil.  When  the  receiver  is  about  half  filled,  stop  the 
distillation  and  seal  the  tube  at  the  constriction,  taking  great 
care  that  the  strength  of  the  walls  is  not  lessened  thereby.  While 
the  tube  is  being  sealed  It  must  be  kept  in  the  freezing  mixture; 
the  tube  and  bath  together  may  be  held  by  a  second  person  at  a 
proper  distance  from  the  blast  lamp.  The  tube  must  be  allowed 
to  cool  while  resting  in  a  perpendicular  position  so  that  the 
liquid  sulphur  dioxide  does  not  come  in  contact  with  the  hot 
parts  of  the  glass. 


FIG.  12. 


SULPHUR  TRIOXIDE.  51 


Critical  Point. 


This  tube  half  filled  with  sulphur  dioxide  is  suited  for  a  demonstration  of 
the  critical  point.  The  critical  temperature  of  sulphur  dioxide  is  155°  C., 
and  the  corresponding  pressure  is  79  atmospheres.  The  following  experiment 
is  taken  from  Nernst's  "Theoretical  Chemistry." 

Insert  the  upper  end  of  the  sealed  tube  containing  the  sul- 
phur dioxide  in  a  cork  stopper  and  clamp  the  whole  at  the 
height  of  25  to  30  cm.  above  the  working  bench,  so  that  the  tube 
makes  an  angle  of  between  30°  and  40°  with  the  horizontal. 
Place  a  small  flame  of  a  Bunsen  burner  under  that  part  of  the 
tube  in  which  the  meniscus  of  the  liquid  is  seen.  The  burner 
should  be  provided  with  a  chimney,  and  the  top  of  the  flame 
should  be  some  distance  below  the  tube.  In  order  to  protect  the 
observer  from  possible  explosion,  place  a  heavy  glass  plate  in 
front  of  the  apparatus.  Observe  that  the  liquid  soon  begins  to 
boil  and  that  its  volume  becomes  smaller  and  smaller;  move 
the  burner  a  little  from  time  to  time  as  it  becomes  necessary. 
When  the  temperature  has  nearly  reached  the  critical  point,  the 
meniscus  becomes  perfectly  flat,  and  appears  as  a  very  fine 
straight  line;  it  disappears  completely  as  soon  as  the  critical 
point  is  reached.  Then  there  is  no  line  of  demarcation  between 
liquid  and  gas.  If  the  flame  is  removed,  a  light  mist  begins  to 
form  almost  immediately  in  the  middle  of  the  tube,  then  sud- 
denly the  meniscus  reappears. 

This  experiment  must  be  carried  out  with  great  caution,  for 
the  pressure  within  the  tube  amounts  to  nearly  80  atmospheres. 
The  apparatus  should  not  be  taken  apart  until  the  tube  and  its 
contents  have  become  perfectly  cold  again. 


28.    Sulphur  Trioxide  by  the  Contact  Process. 

The  fact  that  sulphur  dioxide  will  combine  with  oxygen  when  in  the  presence 
of  finely  divided  platinum  was  known  in  the  first  half  of  the  last  century.  Cl. 
Winkler  showed  as  early  as  1875  that,  by  means  of  such  a  contact-process, 
sulphuric  acid  could  be  made  on  an  industrial  scale  from  mixtures  of  sulphur 
dioxide  and  oxygen.  It  was  not  until  nearly  the  close  of  the  century,  how- 
ever, that  the  doubts  with  regard  to  the  feasibility  of  manufacturing  sulphuric 
anhydride  on  a  large  scale  from  the  gases  evolved  in  the  roasting  of  pyrite 
were  overcome ;  then  the  Badische  Aniline-  und  Sodafabrik  made  public  the 
most  fa-vocable  temperature  for  this  process  and  directed  the  attention  of  a 


52  OXIDES. 

wide  circle  of  chemists  to  the  problem  of  freeing  the  gases  from  catalyzer 
poisons  (cf.  No.  20),  particularly  arsenic  compounds. 

Sulphur  trioxide  exists  in  two  allotropic  modifications:  as  a  mobile  liquid 
(boiling-point  46°)  which  forms  crystals  on  being  sufficiently  cooled  (freezing- 
point  15°),  and  as  a  white  asbestos-like  mass  which  on  warming  volatilizes 
without  previously  melting.  The  latter  is  the  more  stable  modification; 
liquid  sulphur  trioxide  on  standing  goes  over  slowly  of  itself,  or  more  rapidly 
in  the  presence  of  a  trace  of  sulphuric  acid  which  acts  as  catalyzer,  into  this 
asbestos-like  condition. 

The  asbestos-like  form  when  dissolved  in  phosphorus  oxychloride  is  bi- 
molecular,  whereas  the  liquid  sulphur  trioxide  proves  to  be  monomolecular 
when  studied  in  the  same  way.  It  seems  probable,  therefore,  that  the  solid 
form  is  a  polymer  of  the  liquid. 

Make  a  slight  bend  in  a  40  cm,  long  combustion  tube  at  a 
point  about  6  cm.  from  one  end,  and  insert  the  bent  end  in  one 
opening  of  a  two-necked  globular  receiver  (Fig.  16);  to  the  other 
neck  connect  a  glass  tube  leading  to  the  ventilating  flue  of  the 
hood  under  which  the  apparatus  is  constructed.  Make  the 
joints  tight  by  means  of  asbestos  cord.  Close  the  front  end  of 
the  combustion  tube  with  a  cork  through  which  one  arm  of 
a  T-tube  is  inserted  in  order  that  sulphur  dioxide  and  oxygen 
may  be  introduced  at  the  same  time.  Fill  a  section  of  the  com- 
bustion tube,  12  to  15  cm.  long,  with  loosely-packed  platinized 
asbestos  which  is  prepared  by  moistening  the  required  amount 
of  asbestos  with  5  c.c.  of  10%  chlorplatinic  acid  solution,  drying 
and  igniting  the  mass.  After  the  experiment  the  platinized 
asbestos  can  be  purified  by  washing  and  again  igniting,  and  it  is 
then  ready  for  use  again.  The  whole  apparatus  must  be  per- 
fectly dry;  even  the  asbestos  cord  with  which  the  joints  are 
made  tight  must  be  previously  ignited. 

Place  the  combustion  tube  in  an  asbestos  chamber  (cf.  Fig.  4), 
whose  edges  measure  15,  4.5,  and  4.5  cm.  respectively.  There  should 
be  a  wide  slit  in  the  bottom  for  the  entrance  of  the  flame,  and 
an  opening  in  the  cover  to  carry  away  the  combustion  products. 
Place  a  wide  burner  at  some  distance  below  the  combustion 
tube,  and  regulate  the  flame  to  maintain  the  temperature  of 
the  platinized  asbestos  at  about  400°.  This  temperature  may 
be  read  with  a  mercury  thermometer  which  has  been  filled  under 
pressure,  or  it  may  be  estimated  quite  closely  with  a  360° 
thermometer  if  the  latter,  upon  being  placed  inside  the  asbestos 
chamber,  shows  but  a  slow  rise  of  its  thread  above  the  350°  mark. 


OXIDATION   OF   NAPHTHALIN  WITH  SULPHURIC  ACID.        53 

Pass  oxygen  from  a  steel  cylinder  or  a  gasometer  through  a 
wash  bottle  containing  concentrated  sulphuric  acid,  which  serves 
to  dry  the  gas  and  at  the  same  time  to  show  the  speed  with 
which  it  is  being  drawn.  Admit  the  oxygen  through  one  arm  of 
the  T-tube  into  the  combustion  tube,  and  through  the  other 
branch  of  the  T-tube  introduce  sulphur  dioxide  which  is  generated 
by  the  action  of  400  g.  concentrated  sulphuric  acid  upon  100  g. 
copper.  This  gas  must  likewise  be  passed  through  sulphuric  acid, 
and  in  addition  through  a  tube  loosely  filled  with  glass  wool  in  order 
to  free  it  from  the  spray  mechanically  carried  along  from  the 
generating  flask.  Regulate  the  flow  of  the  two  gases  so  that 
a  little  more  oxygen  than  sulphur  dioxide  passes  into  the  con- 
tact tube.  Keep  the  receiver  immersed  in  ice  water;  sulphur 
trioxide  collects  abundantly  either  in  the  liquid  modification  or 
in  the  asbestos-like  form.  The  experiment  takes  about  three  hours. 

For  the  following  experiments  use  the  liquid  form.  In  case 
the  asbestos-like  form  has  been  obtained,  loosen  it  with  a  glass 
stirring  rod  and  place  a  layer  of  it  about  1  cm.  deep  in  each 
of  three  test-tubes;  add  (under  the  hood)  a  drop  of  concentrated 
sulphuric  acid  to  each  tube,  heat  just  to  the  melting-point, 
and  then  allow  to  cool.  To  one  test-tube  add  flowers  of  sul- 
phur from  the  point  of  a  knife  blade;  to  the  second,  powdered 
selenium;  and  to  the  third,  iodine.  In  the  first  test-tube  an 
indigo-blue  solution  is  formed,  in  the  second  a  bluish  green,  and 
in  the  third  likewise  a  bluish  green,  or,  if  considerable  iodine  has 
been  added,  a  brown  solution.  The  compounds,  S2O3  and  SSe03, 
are  produced  in  the  first  and  second  tubes  respectively. 

Oxidation  'of  Naphthalin  with  Sulphuric  Acid. 

The  dissociation  of  sulphur  trioxide,  SO3  =  SO2  +  O,  in  opposition  to  the 
reaction  of  its  synthesis  by  the  contact  process,  is  frequently  utilized  for 
technical  and  analytical  purposes;  by  an  increase  of  temperature,  and  by  the 
presence  of  catalyzers,  such  as  mercury  or  copper  salts,  the  oxidizing  action 
is  considerably  accelerated.  Instead  of  sulphur  trioxide,  fuming  sulphuric 
acid  or  even  ordinary  concentrated  sulphuric  acid  may  be  employed,  although 
the  effect  is  not  then  so  readily  obtained. 

The  most  important  technical  utilization  of  this  oxidizing  power  of  fuming 
sulphuric  acid  is  in  the  transformation  of  naphthalin  into  phthalic  acid,  the 
latter  being  used  in  the  preparation  of  artificial  indigo: 

C10H8      +  9  O  =      CSH0O4       +  H2O  +  2  CO, 
Naphthalin  Phthalic  Acid 


54        .  OXIDES. 

Place  5  g.  naphthalin,  1  g.  mercury,  and  80  g.  concentrated 
sulphuric  acid  in  which  the  remainder  of  the  above  sulphur 
trioxide  has  been  dissolved  (or  80  g.  of  commercial,  fuming  sul- 
phuric acid)  in  a  300  c.c.  retort,  and  heat  the  mixture  slowly, 
almost  to  boiling,  on  a  Babo  funnel.  Insert  the  neck  of  the  retort 
into  a  small  flask,  which  serves  as  a  receiver,  and  cool  the  latter 
with  water.  A  white  sublimate  of  phthalic  anhydride  soon  appears 
in  the  upper  part  and  neck  of  the  retort,  and  the  odor  of  sulphur 
dioxide  becomes  apparent.  From  time  to  time  drive  the  sub- 
limate over  into  the  receiver  by  fanning  the  top  and  neck  of  the 
retort  with  the  flame  of  a  Bunsen  burner. 

—A  When,  at  the  end  of  two  or  three  hours, 

nothing  further  passes  over,  decant  off  the 
pi  ^3  sulphuric  acid  which  has  distilled  into  the 

receiver  and  recrystallize  the  phthalic 

anhydride,  from  75  to  100  c.c.  of  water.  The  filtrate  from  the 
first  crop  yields  more  crystals  on  evaporation. 

Dry  the  phthalic  acid  thus  obtained,  and  distil  it  in  a  test- 
tube  held  nearly  horizontally,  the  closed  end  of  which  is  bent 
downward  at  a  slight  angle,  as  shown  in  Fig.  13.  Water  is  split 
off  from  the  molecule  during  the  sublimation,  and  beautiful 
needles  of  phthalic  anhydride  (melting-point  128°  C.)  are  formed. 
By  another  crystallization  from  water,  2  to  3  g.  of  pure,  per- 
fectly white  phthalic  acid  are  obtained. 

For  the  characterization  and  identification  of  the  phthalic 
acid,  mix  a  little  of  it  with  equal  amounts  of  resorcinol  and 
anhydrous  zinc  chloride.  Heat  this  mixture  slowly  in  a  small, 
dry  test-tube  over  a  small  flame  until  it  sinters  and  then  melts. 
After  heating  a  minute  longer,  cool  the  brownish-red  fusion  and 
dissolve  it  in  a  little  alcohol.  Pour  the  solution  thus  obtained 
into  a  large  beaker  containing  distilled  water,  and  add  a  few  drops 
of  caustic  soda  solution.  A  deep-yellow  solution  is  obtained 
which,  by  reflected  light,  shows  a  beautiful  green  fluorescence 
(synthesis  of  fluorescein). 

29.   Nitrogen  Dioxide. 

Place  200  g.  of  coarse  lumps  of  arsenic  trioxide  in  a  flask,  add 
250  g.  of  concentrated  nitric  acid  (sp.  gr.  1.4),  and  heat  the  mix- 
ture moderately  upon  a  sand  bath  (or  Babo  boiling  funnel). 


NITROGEN  DIOXIDE.  55 

Lead  the  gases  evolved  successively  through  an  empty  wash- 
bottle,  a  U-tube  containing  glass  wool,  a  second  empty  wash- 
bottle,  and  finally  into  a  third  wash-bottle  (surrounded  with  ice) 
in  which  the  oxides  of  nitrogen  are  condensed.  A  mixture  of 
nitrogen  dioxide,  nitrogen  trioxide,1  and  nitric  oxide  is  obtained. 
After  the  evolution  of  gas  from  the  flask  has  ceased,  pass  a  cur- 
rent of  oxygen  through  the  condensed  liquid  in  the  wash  bottle, 
still  keeping  it  surrounded  with  ice,  until  the  color  becomes  a 
pure  yellowish  brown;  only  a  small  amount  of  the  nitrous  gases 
are  lost  during  this  operation.  If  it  is  desired  to  preserve  the 
preparation,  heat  the  flask  containing  the  liquid  nitrogen  dioxide 
cautiously  by  means  of  lukewarm  water  and  distil  the  liquid 
into  a  sealing  tube,  in  exactly  the  same  manner  as  with  sulphur 
dioxide  (p.  50).  Boiling-point  22°  C. 

Cork  stoppers  and  rubber  tubing  are  energetically  attacked 
by  the  oxides  of  nitrogen.  Therefore,  in  fitting  up  the  above 
apparatus,  select  corks  which  fit  tightly  and  protect  them  with 
a  coating  of  vaseline;  where  the  use  of  rubber  connections  is 
unavoidable,  bring  the  ends  of  the  glass  tubing  close  together. 

Gaseous  nitrogen  dioxide  is  of  a  reddish-brown  color  at  ordinary  tem- 
peratures; on  heating  the  color  at  first  becomes  darker  red  on  account  of  a 
progressive  dissociation,  N2O4<=±  2  NO2;  but  on  heating  above  130°  the  color 
again  becomes  lighter  in  consequence  of  a  dissociation  of  the  nitrogen  dioxide 
into  oxygen  and  nitric  oxide :  2  NO2  *=±  2  NO  +  O2. 

Arsenic  acid  may  be  obtained  from  the  residue  in  the  evolution 
flask.  Complete  the  oxidation  by  further  heating  with  con- 
centrated nitric  acid,  and  evaporate  the  solution  to  a  fairly  thick 
sirup.  If  exactly  the  right  concentration  of  the  sirupy  solution 
is  obtained,  it  will  form  a  nearly  solid  mass  of  crystals  of  arsenic 
acid  on  standing  in  the  ice  chest. 

30.    Chromic  Oxide  in  the  Dry  Way  from  a  Chromate. 

The  preparation  of  chromic  oxide  by  the  reduction  of  an  aqueous  solution 
of  a  chromate  to  one  of  a  chromic  salt,  with  subsequent  precipitation  and 
ignition  of  chromic  hydroxide,  is  not  a  convenient  process  because  of  the 
difficulty  and  tediousness  of  filtering  and  washing  the  very  voluminous  pre- 
cipitate. The  method  originated  by  Wohler,  1827,  which  is  that  given  in  the 

1  Nitrogen  trioxide  N2O3  can  exist  only  in  the  liquid  condition.  It  is 
unstable,  and  when  vaporized  dissociates  into  NO  and  NO2  (or  N2O4). 
v.Wittorff,  Z.  Anorg.  Chem.  41,  85  (1904). 


56  HYDRIDES 

following  directions,  is,  therefore,  preferable.  The  chromate  is  heated  in  a 
crucible  with  ammonium  chloride,  whereby  the  ammonium  radical  is  oxidized 
to  water  and  nitrogen  while  the  chlorine  combines  with  the  alkali  metal  of  the 
chromate. 

Mix  147  g.  potassium  pyrochromate  (one-half  mol.)  intimately 
with  an  equal  weight  of  ammonium  chloride,  and  heat  the  mixture 
in  a  clay  crucible,  in  a  charcoal  furnace,  until  no  more  vapors 
are  given  off.  After  cooling,  boil  the  brittle  contents  of  the 
crucible  repeatedly  with  fresh  portions  of  water  until  all  the 
soluble  salt  has  been  removed.  Compute  the  yield  of  the  dried 
preparation  in  percentage  of  the  theoretical.  Dependent  prep- 
aration, Chromium  No.  3. 

31.    Cuprous  Oxide  from  Fehling's  Solution. 

Dissolve  50  g.  crystallized  cupric  sulphate,  75  g.  potassium- 
sodium  tartrate,  and  75  g.  sodium  hydroxide  in  a  porcelain  evap- 
orating dish  by  warming  slightly  with  600  c.c.  of  water.  To  the 
blue  solution  thus  obtained  add  100  g.  cane-sugar  and  heat  to 
boiling,  whereby  the  blue  color  gradually  disappears  and  a  heavy, 
dark-red  precipitate  of  cuprous  oxide  is  formed.  Free  the  pre- 
cipitate as  completely  as  possible  from  alkali  by  pouring  off  the 
solution  and  washing  repeatedly  by  decantation  with  water. 
Then  bring  the  precipitate  itself  upon  a  hardened  filter  and  wash 
it  with  considerable  water,  and  at  last  with  a  little  alcohol.  Dry 
the  cuprous  oxide  in  the.  hot  closet.  Yield,  14  g. 

HYDRIDES. 

A  few  metals  combine  directly  with  hydrogen  at  definite  temperatures  to 
form  hydrides  (e.g.,  cerium  hydride  No.  32). 

Copper  hydride  is  formed  in  aqueous  solution  by  treating  cupric  salts  with 
very  strong  reducing  agents  (No.  33). 

Certain  products  which  are  formed  by  the  occlusion  of  hydrogen  on  the 
part  of  some  of  the  heavy  metals,  and  are  to  be  regarded  as  solid  solutions, 
should  not  be  confused  with  the  true  hydrides;  e.g.,  palladium-hydrogen. 

32.    Cerium  Hydride. 

If  cerium  dioxide  is  reduced  by  means  of  metallic  magnesium  in  an  atmos- 
phere of  hydrogen,  the  metallic  cerium,  as  fast  as  it  is  set  free,  combines  with 
the  gas  to  form  cerium  hydride. 


CERIUM  HYDRIDE.  57 

Ignite  a  few  grams  of  eerie  ammonium  nitrate  (No.  171),  or 
cerous  ammonium  nitrate,  at  first  gently,  and  finally  with  the  blast 
lamp,  and  mix  the  pure,  yellow  cerium  dioxide  thus  obtained 
with  powdered  magnesium  in  the  proportion  of  172  parts  CeO2 
to  64  parts  Mg.  Place  the  mixture  in  a  boat  and  introduce  it 
into  a  short  combustion  tube,  one  end  of  which  is  connected 
through  a  sulphuric  acid  wash-bottle  with  a  hydrogen  generator, 
while  the  other  end  can  be  closed  when  required.  After  the 
air  in  the  apparatus  has  been  completely  replaced  by  hydrogen, 
heat  the  tube  in  a  short  combustion  furnace,  at  first  gently,  until 
all  moisture  is  removed;  then  close  the  exit  end  of  the  combus- 
tion tube  and  open  the  cock  of  the  generator  wide  so  that  the 
reaction  mixture  stands  under  a  slight  pressure  of  hydrogen. 
On  now  heating  the  tube  with  the  full  flame,  the  mixture  glows, 
the  walls  of  the  tube  above  it  become  blackened,  and  gas  bubbles 
pass  through  the  wash  bottle,  at  first  rapidly  and  then  less  and 
less  frequently,  corresponding  to  the  rate  of  consumption  of  the 
hydrogen.  Heat  the  mixture  five  or  ten  minutes  longer,  and 
then  allow  it  to  cool  under  hydrogen  pressure.  The  reddish-brown, 
fairly  compact  reaction-product  consists  of  a  mixture  of  cerium 
hydride  and  magnesium  oxide.  Break  pieces  of  it  from  the  boat 
and  set  fire  to  them  with  a  match;  the  material  burns  with  a 
hydrogen  flame  to  a  nearly  white  ash.  While  the  hydrogen  is 
burning,  occasional  flashes  occur  from  the  ignition  of  particles  of 
unoxidized  magnesium.  It  is  possible  to  keep  cerium  hydride 
for  a  long  time  in  sealed  vessels. 

By  heating  calcium  turnings  to  a  dull  red  heat  in  an  atmos- 
phere of  hydrogen,  calcium  hydride,  CaH2,  may  be  prepared  in 
an  analogous  manner. 

33.    Copper  Hydride. 

From  100  c.c.  of  a  5%  solution  of  barium  hypophosphite 
(No.  96),  precipitate  all  of  the  barium  by  the  addition  of  about 
18  c.c.  of  2-normal  sulphuric  acid,  and  treat  the  clear  filtrate 
at  the  room  temperature  with  four  grams  of  copper  sulphate 
dissolved  in  15  c.c.  of  water.  After  some  time,  or  more  quickly 
if  heated  to  30°,  the  solution  becomes  green,  and  dark-brown 
copper  hydride  is  precipitated  which  to  some  extent  adheres  to 
the  sides  of  the  glass  vessel  as  an  iridescent  film.  Filter  off  the 


58  ACIDS,  BASES,  AND  SALTS. 

precipitate,  wash  it  with  water,  and  dry  the  product  in  a  vacuum- 
desiccator. 

Heat  a  portion  of  the  copper  hydride  in  a  small  test-tube.  It 
decomposes  suddenly  into  red  copper  and  hydrogen;  the  latter 
may  be  ignited  at  the  mouth  of  the  tube. 

Cover  a  second  portion  with  concentrated  hydrochloric  acid; 
hydrogen  is  evolved,  which,  in  consequence  of  the  admixed  spray 
of  copper  compounds,  burns  with  a  deep-blue  flame.  From  the 
solution  in  the  test-tube,  the  addition  of  a  little  water  precipitates 
white  cuprous  chloride. 

ACIDS,    BASES,    AND    SALIS. 

The  majority  of  the  simple  compounds  are  electrolytes,  that  is,  they  are 
acids,  bases,  or  salts.  The  preparation  of  simple  electrolytes  can  take  place 
according  to  the  methods  outlined  on  p.  44  as  generally  applicable  for 
simple  compounds;  as  peculiar  to  electrolytes,  the  formation  of  a  compound 
by  the  association  of  its  tons,  presents  a  special  case  of  synthesis  from  the 
elements.  In  the  preparation  of  pure  electrolytes  in  solution  by  bringing 
together  the  necessary  ions,  it  is  essential  that  foreign  ions  should  be  removed. 
Thus  soluble  hydroxides  may  be  prepared  by  precipitating  solutions  of  the 
corresponding  sulphates  with  an  equivalent  amount  of  barium  hydroxide, 
the  barium  and  sulphate  ions  being  removed  as  insoluble  barium  sulphate, 
and  only  the  desired  hydroxide  and  its  ions  remaining  in  solution  (cf.  Nos.  36 
and  124).  In  a  similar  way  acids  may  be  obtained  by  the  interaction  of 
barium  salts  and  sulphuric  acid  (No.  33).  Potassium  iodide,  likewise,  may 
be  prepared  conveniently  by  the  double  decomposition  of  ferrous  iodide  with 
potassium  carbonate,  whereby  iron  and  carbonate  ions  are  precipitated  in 
the  form  of  an  insoluble  compound  (No.  39). 

ELECTROLYTIC  DISSOCIATION  is  distinguished  from  simple  dissociation  by 
the  fact  that  the  products  of  dissociation  are  electrically  charged.  Further- 
more, the  extent  of  electrolytic  dissociation  depends  in  the  highest  degree  upon 
the  nature  of  the  solvent  medium. 

2  HI  =  H2  +  I2       (non-electrolytic  dissociation). 

HI  =  H++  I~     (electrolytic  dissociation). 

This  double  possibility  of  dissociation  enables  substances  to  react  in  different 
ways.  According  to  the  first  of  the  above  equations,  hydrogen  iodide  is  a 
reducing  agent;  according  to  the  second,  it  is  an  acid. 

Electrolytic  dissociation  can  be  detected  and  measured  by  physical  methods, 
either  by  estimating  the  size  of  the  molecule  of  the  dissolved  substance  by 
one  of  the  osmotic  methods,  or  by  determining  the  conductivity  of  the  solution 
(Arrhenius) ;  cf .  No.  34.  As  Nernst  has  suggested,  the  process  of  electrolytic 
dissociation  may  be  formulated  if  electricity  is  regarded  as  something  material, 
and  its  elementary  quantities  (the  electrons)  are,  in  much  the  same  way  as 
the  atoms,  represented  by  the  special  symbols  ©  and  0,  whereby  the  question 


ACIDS,  BASES,  AND   SALTS.  59 

as  to  whether  or  not  positive  electricity  is  merely  the  absence  of  negative 
electricity  is  left  entirely  undecided.  The  process  of  electrolytic  dissociation 
can  then  be  considered  as  a  double  decomposition  ;  the  elementary  atoms  con- 
tained in  the  compound  unite  with  the  neutral  electricity  which  is  present 
everywhere  (neutron)  in  such  a  manner  that  a  compound  of  one  of  the  atoms 
with  a  positive  electron  (cation)  and  another  compound  of  the  other  atom 
with  a  negative  electron  (anion)  result: 


AB  +    ©© 

binary  compound         neutron          cation        anion 

LAW  OF  DILUTION.     The  application  of  the  law  of  mass-action  to  the 
dissociation  of  a  binary  electrolyte  gives  the  following  expression: 

_  [cation]  [anion]  _ 

/c. 


[undissociated  compound] 

Thus,  in  a  given  case,  if  the  fraction  of  a  gram  molecule  of  a  binary  electrolyte 
which  has  undergone  dissociation  is  denoted  by  a  (the  degree  of  dissociation), 
then  1  —  a  is  the  undissociated  fraction  of  the  gram  molecule  which  is  in 
equilibrium  with  the  ions,  and  at  the  dilution  v  the  concentration  of  the 

undissociated  compound  is and  that  of  each  ion  is—.    Then  the  mass- 

v  v 

action  law  shows  the  following  relation: 


k,        or      — — •  =  k.    (Ostwald's  Law  of  Dilution.) 


The  maximum  possible  value  for  a,  the  degree  of  dissociation,  is  1,  and  this 
represents  complete  dissociation;  it  is  evident  that  this  is  approached  in  pro- 
portion as  v  is  made  large.  Experiment  37  illustrates  a  case  of  dissociation 
with  progressive  dilution. 

IONIZATION  TENDENCY,  OR  ELECTROAFFINITY.  A  measure  of  the  force 
with  which  the  electron  is  held  to  the  material  atom  of  an  ion  is  given  by  the 
potential  which  is  necessary  to  effect  the  discharge  of  the  ion  in  an  electrolysis 
(decomposition  potential).  This  force  is  variously  known  as  ionization 
tendency  or  electroa ffinity .  In  the  following  table,  the  tension  in  volts  is  given 
which  is  necessary  for  the  discharge  of  a  few  of  the  metal  ions  from  their 
normal  solutions,  on  the  basis  of  the  discharge-potential  of  hydrogen  being 
taken  arbitrarily  as  zero.1 

Mg  +  1.482  Cd  +  0.420  Pb  4- 0.151 

Al    +  1.276  Fe  +  0.344  Cu  -  0.329 

Mn  +  1.075  Co  +  0.232  Ag  -  0.753 

Zn  +  0.770  Ni  +  0.228 


1  The  choice  of  an  arbitrary  zero-point  is  necessary,  since  it  is  only  differ- 
ences in  potential  which  can  actually  be  measured. 


60  ACIDS    BASES    AND    SALTS. 

In  this  potential  series  the  distinction  which  has  always  been  recognized 
between  noble  and  base  metals  is  again  expressed.  Elements  with  a  high 
ionizing  tendency  possess  to  a  marked  degree  the  ability  to  form  simple  ions. 
Simple  ions  of  the  noble  metals  are  less  stable,  so  that  the  number  and  impor- 
tance of  the  simple  salts  of  these  metals  is  much  less  than  that  of  their  com- 
plex salts.  (Cf.  Chapters  IV  and  V.) 

HYDROLYTIC  DISSOCIATION,  OR  HYDROLYSIS.  The  formation  of  a  salt  by 
the  neutralization  of  an  acid  and  a  base  is  a  reversible  process.  If  M  denotes 
a  metal  and  R  an  acid  radical,  then  the  reaction  of  equilibrium  is: 

Neutralization 

MOH  +  HR    +±    MR  +  H2O 
Hydrolysis 

If  the  equation  is  read  from  right  to  left  it  represents  an  hydrolysis,  i.e.,  the 
breaking  up  of  a  salt  into  an  acid  and  a  base.  Neutralization  and  hydrolysis 
represent,  therefore,  reciprocal  processes.  It  is  important  in  the  preparation 
of  salts  to  know  which  of  these  reactions  preponderates.  The  mass-action 
law  applied  to  this  reversible  reaction  gives: 

[HR]  [MOH]  _ 

[MR]  [HOH] 

Salts,  acids,  bas?s,  and  water  are  dissociated  electrolytically  according  to  the 
following  equations: 

1.  MR  =  M+  +  R~ 

2.  HR  =  H+  +  R~ 

3.  MOH  =  M+  +  OH~ 

4.  HOH  =  H+  +  OH~ 

The  corresponding  dissociation  constants  Klt  K2,  K3)  K4,  may  be  given  as 
follows  :  1 

[M+]  [R~]   .  [H+]  [R-]  .  [M*]  [OH"]  [H+T  [QH~] 


[MR]  [HR]  [MOH]  [HOH] 

If  in  the  above  equation  of  hydrolysis  the  values  for  the  concentration  of 
undissociated  MR,  HR,  MOH,  and  HOH  obtained  respectively  from  the  last 
four  expressions  are  inserted,  the  equation 

~        K-   -  K 


is  obtained.  This  equation  expresses  all  the  essential  phenomena  in  neu- 
tralization and  hydrolysis.  It  shows  that  hydrolysis  will  be  greater  if  the 
value  of  K4  becomes  larger;  or  in  other  words  if  the  temperature  is  raised,  since 


1  The  law  of  mass-action,  as  is  well  known,  does  not  apply  rigidly  to  the 
dissociation  of  strong  electrolytes.  It  may  serve,  however,  as  a  qualitative 
guide  in  estimating  the  dissociation  of  salts  and  of  strong  acids  and  bases. 
(Translators.) 


DETECTION  OF  ELECTROLYTIC  DISSOCIATION.  61 

the  dissociation  of  water,  although  extremely  small,  increases  quite  rapidly 
with  rise  of  temperature.  The  hydrolysis  is  also  greater  as  K2  and  K3  are 
smaller,  or  in  other  words  as  the  acid  and  base  in  question  are  weaker.  Inas- 
much as  the  weakness  of  the  base  usually  increases  with  the  valence  of  the 
metal,  it  follows  that  salts  of  the  trivalent  metals  (Fe,  Al,  Sb,  Bi)  are  hydrolyzed 
to  a  greater  extent  than  the  salts  of  metals  having  a  lower  valence  (Mn,  Ba, 
etc.)  (cf.  Nos.  40,  41).  The  halides  of  the  tetravalent  metals  (Snj  Si, 
Ti)  are  hydrolyzed  to  an  especially  marked  extent  (cf.  Nos.  50  to  52), 
and  it  is,  therefore,  possible  to  prepare  such  salts  pure  only  when  water  is 
completely  excluded.  If  sulphur  chloride  is  regarded  as  the  sulphur  salt  of 
hydrochloric  acid  (cf.  No.  45),  this  compound  then  forms  the  extreme  of  the 
hydrolyzable  chlorides,  inasmuch  as  the  metallic  character  of  the  element 
combined  with  chlorine  has  entirely  disappeared. 

Salts  of  weak  acids  and  weak  bases,  such  as,  for  example,  the  sulphides  of 
aluminium  (Nos.  4  and  5)  and  titanium,  can  only  be  prepared  when  out  of 
contact  with  water.  After  they  have  been  prepared  synthetically  in  a  crys- 
tallized form,  they  can  be  kept  for  some  time  without  change  on  account  of  the 
fact  that  the  velocity  at  which  compact  substances  enter  into  reaction  is  often 
very  slight.  (Cf.  TiS2,  No.  58.) 

If  the  phosphides,  nitrides,  and  certain  carbides  are  regarded  as  the  metal 
salts  of  phosphine,  ammonia,  and  the  hydrocarbons,  that  is  to  say,  as  the  salts 
of  hydrogen  compounds  the  acid  nature  of  which  is  almost  infinitesimal  (the 
acidic  constants,  K2,  of  which  are  therefore  extremely  small),  it  then  becomes 
easy  to  understand  why  these  substances  can  be  prepared  only  in  the  absence 
of  water  (best  by  direct  union  of  the  elements  themselves).  The  readiness 
with  which  the  preparation  of  phosphine  from  calcium  or  magnesium  phos- 
phide (No,  63),  of  acetylene  from  calcium  carbide  (No.  64),  and  of  ammonia 
from  magnesium  or  boron  nitride  (Nos.  60,  61),  is  carried  out,  depends  on  the 
ease  with  which  the  salts  of  such  extremely  weak  acids  hydrolyze. 


(a)  Acids  and  Bases. 
34.   Physico-chemical  Detection  of  Electrolytic  Dissociation. 

To  demonstrate  the  dissociation  of  a  substance  in  aqueous  solu- 
tion, determine  the  molecular  weight  of  nitric  acid  by  the  freezing- 
point  method,  first  with  nitrobenzene  and  then  with  water  as  the 
solvent.  Dehydrate  some  concentrated  nitric  acid  by  distilling 
it  with  an  equal  volume  of  concentrated  sulphuric  acid,  and 
free  the  distillate  from  oxides  of  nitrogen  by  passing  dry  air 
through  it. 

Determine  the  molecular  weight  by  measuring  the  lowering  of 
the  freezing-point  of  nitrobenzene,  using  about  25  g.  of  the  sol- 
vent, and  0.2,  0.4,  1.0,  and  1.5  g.  of  nitric  acid,  then  repeat,  using 


62  ACIDS   AND  BASES. 

about  22  g.  of  water  instead  of  nitrobenzene.  The  details  for 
carrying  out  the  determinations  may  be  found  in  H.  Biltz;  Prac- 
tical Methods  for  Determining  Molecular  Weights,  Translated  by 
H.  C.  Jones,  1899.  In  nitrobenzene  the  molecular  weight  corre- 
sponds approximately  to  the  formula  HNO3,  whereas  in  water  the 
molecular  weight  is  about  half  as  large. 

In  order  to  show  that  this  dissociation  causes  the  aqueous  solu- 
tion to  become  a  conductor  of  electricity,  insert  two  platinum 
electrodes  in  a  beaker  of  water,  and  connect  them  through 
an  electric  incandescent  lamp  with  the  terminals  of  a  lighting 
circuit.  The  lamp  does  not  glow  because  water  is  a  very  poor 
conductor  of  electricity,  and  the  circuit  is  therefore  practically 
open.  As  soon,  however,  as  a  few  drops  of  nitric  acid  are  added 
to  the  water,  the  lamp  begins  to  glow,  and  at  the  same  time  an 
evolution  of  gas  takes  place  at  the  electrodes. 

Test  the  conductivity  of  nitrobenzene  after  the  same  manner, 
first  alone,  and  then  with  the  addition  of  a  little  anhydrous  nitric 
acid;  the  lamp  does  not  glow  in  either  case. 

35.   Hydrobromlc  Acid. 

(a)  From  the  Elements.  Connect  in  series,  a  Kipp  hydrogen 
generator,  a  wash  bottle  containing  sulphuric  acid,  a  wash  bottle 
containing  80  g.  of  bromine,  and  a  glass  tube  about  25  cm.  long 
and  1  cm.  wide  which  is  filled  for  a  space  of  8  to  12  cm.  with 
loosely-packed,  platinized  asbestos.  (Cf.  No.  28.)  Place  a 
Bunsen  lamp  with  a  flame  spreader  under  the  tube.  Connect  the 
farther  end  of  the  contact  tube  with  a  U-tube  containing  glass 
beads  and  moist  red  phosphorus,  then  with  a  *wash  bottle  made 
from  a  test  tube  which  contains  1  c.c.  of  water,  and  finally  with 
two  wash  bottles  each  containing  35  c.c.  of  water.  Keep  the 
latter  cooled  during  the  experiment,  first  with  ice-water  and  later 
with  a  mixture  of  salt  and  ice. 

At  the  beginning  of  the  experiment  disconnect  the  train  at  a 
point  between  the  contact  tube  and  the  phosphorus  tube.  Fill 
the  apparatus  up  to  this  point  with  hydrogen,  then  heat  the  con- 
tact layer  to  faint  redness,  and  when  the  gas  escaping  from  this 
tube  becomes  colorless,  reconnect  the  rest  of  the  apparatus. 
Regulate  the  amount  of  bromine  vapor  by  pouring  warm  (not 


HYDROBROMIC    ACID.  63 

hot)  water  from  time  to  time  into  the  beaker  which  surrounds 
the  bromine  bottle.  The  gases  after  passing  the  contact  layer 
must  contain  no  uncombined  bromine,  and  must  therefore  be 
colorless. 

From  the  first  receiver  fuming  hydrobromic  acid  is  obtained, 
and  the  yield  can  be  determined  by  weighing  the  flask  before  and 
after  the  experiment. 

(6)  By  the  Intermediate  Formation  of  Phosphorus  Bromide. 
Provide  a  500  c.c.  flask  with  a  stopper  and  dropping  funnel  and 
connect  it  by  means  of  delivery  tubing  with  two  U-tubes  placed 
in  series.  From  the  last  U-tube  carry  a  bent  delivery  tube 
through  a  tight-fitting  cork  well  towards  the  bottom  of  a  flask 
which  is  to  serve  as  a  receiving  vessel.  Provide  this  flask  with  an 
exit  tube  starting  from  just  inside  the  cork  and  bending  down- 
ward externally  so  as  to  nearly  reach  the  surface  of  some  water 
placed  in  another  flask.  Do  not,  during  the  process  which  follows, 
allow  the  tube  entering  either  of  the  last-mentioned  flasks  to  dip 
beneath  the  surface  of  the  solution  contained  in  them. 

Place  in  the  evolution  flask,  first  a  layer  of  25  g.  sand,  then  upon 
this  a  mixture  of  100  g.  sand  and  25  g.  red  phosphorus,  and  moisten 
the  whole  with  45  c.c.  water.  Fill  the  first  U-tube  with  glass 
beads,  the  second  with  glass  beads  mixed  with  moist  red  phos- 
phorus, and  in  the  receiving  flask  place  80  c.c.  of  water.  The 
delivery  tube  which  enters  the  receiver  must  reach  only  to  a 
point  just  above  the  surface  of  the  liquid,  since  if  it  dipped  into 
the  solution,  the  latter  might  be  sucked  back  into  the  evolution 
flask. 

Wrap  the  evolution  flask  in  a  towel,  because  it  sometimes  breaks 
at  the  beginning  of  the  experiment;  cool  the  first  U-tube  with  a 
mixture  of  salt  and  ice,  and  the  receiver  with  ice.  Introduce 
from  the  funnel  200  g.  bromine  drop  by  drop  into  the  mixture  of 
sand  and  phosphorus.  At  the  beginning  of  the  reaction  cool  the 
evolution  flask  by  placing  it  in  a  dish  containing  cold  water. 
Yield,  about  250  g.  of  concentrated  hydrobromic  acid,  which  is 
collected  in  the  first  receiving  flask.  Determine  its  density  by 
means  of  a  Westphal  balance.  Dependent  preparations:  Cupric 
Bromide  (No.  37),  Ammonium  Tribromide  (No.  69),  Hexammine- 
nickelous  Bromide  (No.  128),  Praseocobalt  Bromide  (No.  135). 
Hydriodic  acid  may  be  prepared  by  a  similar  method. 


64  ACIDS    AND    BASES. 

36.   Thallous  and  Thallic  Hydroxides. 

Mix  13'.3  g.  of  thallous  nitrate1  (1/20  mol.)  and  3  g.  concentrated 
sulphuric  acid  in  a  platinum  crucible,  and  drive  off  the  volatile 
acid  slowly  by  heating  the  upper  edge  of  the  crucible  with  a 
Bunsen  burner  which  is  held  in  the  hand  and  moved  about  so  as 
to  avoid  spattering.  The  decomposition  is  complete  when  the 
contents  of  the  crucible  no  longer  give  off  acid  vapors.  Dissolve 
the  thallous  sulphate  thus  prepared  in  50  c.c.  of  water,  and  pre- 
cipitate barium  sulphate  from  the  boiling  solution  by  adding  a 
hot  solution  of  8  g.  crystallized  barium  hydroxide,  the  amount 
of  the  latter  reagent  being  slightly  in  excess  of  that  theoreti- 
cally required.  After  filtering,  precipitate  the  excess  of  barium 
hydroxide  by  the  careful  addition  of  very  dilute  sulphuric  acid. 
A  drop  of  the  solution  should  not  at  the  last  give  a  precipitate 
either  with  barium  hydroxide  or  with  sulphuric  acid.  Evaporate 
the  solution  to  a  volume  of  about  25  c.c.,  filter  again  and  then 
evaporate  to  about  5  c.c.,  after  which,  place  the  evaporating  dish 
in  a  desiccator  over  dry  lime.  Suck  the  crystals  of  thallous 
hydroxide  free  from  liquid  in  a  small  filter  funnel  and  obtain  a 
second  crop  by  further  evaporation  of  the  mother  liquor.  Neu- 
tralize the  last  mother  liquor  with  nitric  acid,  and  use  the 
resulting  solution  of  thallous  nitrate  for  the  reaction  described 
below. 

Thallous  hydroxide  forms  light-yellow  crystals  which  are 
readily  soluble  in  water;  the  aqueous  solution  of  these  crystals  is 
strongly  alkaline,  its  electrical  conductivity  being  nearly  the  same 
as  that  of  a  sodium  hydroxide  solution  of  the  same  concentration. 
Filter  paper  moistened  with  thallous  hydroxide  serves  as  a 
reagent  for  detecting  the  presence  of  ozone ;  brown  thallic  oxide  is 
formed. 

Reactions  of  Thallous  Salts.  Solutions  of  thallous  salts  when 
treated  with  iodide  solutions  give  a  yellow,  very  difficultly  soluble 
precipitate  of  thallous  iodide  (analogous  to  Agl);  with  chloro- 


1  In  order  to  obtain  thallous  nitrate  from  metallic  thallium,  the  latter  is 
cut  into  small  pieces  and  dissolved  in  the  calculated  amount  of  dilute  nitric 
acid.  The  hot,  slightly  acid  solution  is  freed  from  lead  by  means  of  hydrogen 
sulphide,  and  the  filtrate  is  evaporated  until  the  nitrate  crystallizes  out. 
The  crystals  are  sucked  dry,  and  the  mother  liquor  together  with  the  wash- 
water  is  evaporated  for  more  crystals.  Yield,  theoretical. 


CUPRIC  AND  CUPROUS  BROMIDES.  65 

platinic  acid,  they  give  light-yellow,  difficultly-soluble  thallium 
chloroplatinate,  T12  [PtCl6]  (analogous  to  K2  [PtClJ) ;  with  ammo- 
nium sulphide,  or  hydrogen  sulphide  together  with  sodium  acetate, 
they  give  a  deep-brown  precipitate  of  thallous  sulphide  (analogous 
to  Ag2S). 

Hydrated  Thallic  Oxide  (analogous  to  hydrated  ferric  oxide). 
Oxidize  a  part  of  the  thallous  nitrate  solution  by  the  addition  of 
an  excess  of  bromine,  and  precipitate  the  resulting  brown  solution 
by  adding  concentrated  ammonia.  The  very  fine  precipitate 
settles  slowly.  Wash  it  by  decantation  with  water  containing 
ammonia  until  the  decanted  liquid  is  free  from  bromine,  and 
then  drain  it  upon  a  filter.  Dry  the  product  in  a  desiccator  over 
sulphuric  acid,  for  at  higher  temperature  it  loses  water  and  becomes 
thallic  oxide. 

(b)    Halogen  Compounds. 
37.    Cupric  and  Cuprous  Bromides. 

Dissolve  16  g.  of  cupric  oxide  in  a  solution  of  hydrobromic  acid 
containing  33  g.  of  pure  HBr.  Filter  the  resulting  solution  and 
evaporate  it  to  a  small  volume,  whereby  the  color  becomes  very 
dark.  Place  the  evaporating  dish  with  its  contents  in  a  desic- 
cator over  sulphuric  acid,  preferably  in  a  vacuum;  an  almost 
solid  mass  of  crystals  forms  in  a  few  days.  Break  up  this  mass 
thoroughly  once  each  day  to  accelerate  the  drying  process. 
Cupric  bromide  is  of  a  deep-black  color  and  glistens  somewhat 
like  iodine. 

Heat  a  portion  of  the  cupric  bromide  cautiously  in  an  evapo- 
rating dish;  it  loses  bromine  and  is  changed  into  white  cuprous 
bromide.  Cuprous  bromide  is  practically  insoluble  in  water. 

The  color  of  the  cupric  bromide  solution  varies  with  the  dilution; 
the  most  dilute  solutions,  which  contain  practically  all  the  copper 
in  the  ionic  state,  are  light  blue;  the  most  concentrated  ones  have 
the  peculiar,  dark-brown  color  of  undissociated  cupric  bromide; 
other  concentrations  have  intermediate  shades.1 


1  The  light-blue  color  of  the  dilute  solutions  is  the  characteristic  color  of 
all  solutions  of  cupric  ions.  The  changes  in  color  with  increasing  concentra- 
tion may  be  due,  certainly  in  part,  to  an  increasing  proportion  of  undissociated 
molecules,  but  it  is  also  quite  certain  that  changes  in  the  state  of  hydration  of 
the  dissolved  salt  have  a  large  influence  in  altering  the  color.  (Translators.) 


66  HALOGEN    COMPOUNDS. 

38.  Cuprous  Chloride. 

Treat  50  g.  crystallized  cupric  sulphate  and  25  g.  sodium  chloride 
(or  37  g.  crystallized  cupric  chloride)  in  a  flask  with  150  g.  concen- 
trated hydrochloric  acid  and  20  g.  copper  turnings.  Heat  upon 
the  water  bath  until,  at  the  end  of  about  an  hour,  the  green  color 
has  disappeared.  Pour  the  clear  solution  into  a  liter  of  water 
containing  a  little  sulphurous  acid,  whereby  white  cuprous  chloride, 
insoluble  in  dilute  hydrochloric  acid,  is  precipitated.  Wash  the 
precipitate  by  decantation  with  water  containing  sulphurous 
acid,  finally  drain  it  with  suction  and  wash  successively  with 
glacial  acetic  acid,  alcohol,  and  ether.  Dry  it  in  the  hot  closet. 
Yield,  about  20  g. 

39.  Potassium  Iodide. 

By  the  action  of  iodine  upon  a  caustic  potash  solution,  a  mixture  of  potas- 
sium iodide  and  iodate  is  produced,  and  this  can  be  reduced  completely  to  the 
iodide  by  heating  with  charcoal.  By  extracting  the  mass  with  water  and 
recrystallizing,  the  product  can  be  purified.1  Potassium  iodide  may  also  be 
prepared  without  the  intermediate  formation  of  iodate,  by  the  interaction  of 
ferrous  iodide  (which  can  be  obtained  synthetically  from  the  elements)  and 
potassium  carbonate. 

Shake  7  or  8  g.  iron  filings  and  50  c.c.  water  in  an  Erlenmeyer 
flask  with  25  g.  iodine  added  in  small  portions.  Warm  the  mix- 
ture somewhat  until  all  of  the  iodine  has  combined,  and  the  color 
of  the  solution  has  become  deep  yellow  (ferrous  iodide) ;  then  pour 
off  the  liquid  from  the  excess  of  iron.  Add  five  grams  more  of 
iodine  to  the  solution  and  heat  until  it  is  dissolved.2  Pour  this 
solution  into  a  boiling  solution  of  17  g.  potassium  carbonate  in 
50  c.c.  water.  The  mixture,  which  at  first  is  very  thick,  becomes 
more  fluid  upon  further  heating,  since  the  precipitate  assumes  a 
more  compact  form.  A  little  of  the  solution  when  filtered  must 
be  perfectly  colorless,  and  free  from  iron;  if  this  is  not  the  case, 
add  a  little  more  potassium  carbonate  to  the  boiling  solution. 
Evaporate  the  filtrate  to  a  small  volume  in  a  porcelain  dish,  filter 
again,  and  evaporate  further  in  a  beaker  until  crystals  begin  to 


1  Cf.  Preparation  of  potassium  bromate  and  bromide,  No.  78. 

2  The  extra  addition  of  the  iodine  serves  to  partially  oxidize  the  ferrous 
salt,  and  the  subsequent  precipitate  thus  contains  hydrated  ferric  oxide  and 
is  readily  filtered ;  ferrous  carbonate  would  be  very  difficult  to  filter. 


BARIUM  CHLORIDE  FROM  WITHERITE.  67 

separate.  Then  allow  the  solution  to  evaporate  slowly  by  placing 
the  beaker  in  a  warm  place  (as  on  top  of  the  hot  closet).  Drain 
the  crystals  in  a  funnel,  wash  them  with  a  little  cold  water,  and 
save  the  mother  liquor  for  another  crop  of  crystals.  Yield,  25 
to  35  g. 

The  potassium  iodide  when  dissolved  in  a  little  water  and  acidi- 
fied should  not  show  any  yellow  color  (free  iodine),  which  would 
indicate  the  presence  of  iodate  in  the  salt. 

40.   Barium  Chloride  from  Witherite. 

When  a  sample  of  witherite  is  dissolved  in  hydrochloric  acid,  the  resulting 
solution  of  barium  chloride  contains  iron  as  well  as  other  impurities.  The 
simplest  way  to  remove  the  iron  is  to  oxidize  it  with  a  little  chlorine  water, 
and  then  precipitate  it  by  adding  an  excess  of  the  powdered  mineral  (cf.  the 
barium  carbonate  method  for  the  analytical  separation  of  metals,  the  salts 
of  which  hydrolyze  to  different  degrees;  see  p.  61).  The  more  difficultly 
soluble  barium  chloride  is  freed  from  any  calcium  and  strontium  chlorides  by 
crystallization. 

First  determine  with  a  Westphal  balance  the  specific  gravity 
of  the  pure,  concentrated  hydrochloric  acid  which  is  to  be  used, 
find  the  percentage  of  HC1  by  referring  to  specific  gravity  tables, 
and  compute  the  quantity  of  the  solution  required  to  dissolve 
100  g.  of  witherite.  Place  this  amount  in  a  two-liter  flask,  dilute 
it  with  water  to  a  volume  of  1500  c.c.,  add  the  powdered  witherite 
and  heat  until  the  mineral  is  dissolved.  Impurities  such  as  silicates 
may  remain  as  insoluble  residue.  Add  50  c.c.  of  chlorine  water, 
then  5  to  10  g.  more  of  witherite,  and  allow  the  solution  to  stand 
in  a  warm  place  with  frequent  shaking.  Next  morning  again  add 
from  2  to  5  g.  of  witherite.  From  time  to  time  filter  a  little  of  the 
solution  and  test  with  potassium  thiocyanate  for  iron.  When  all 
the  iron  has  been  precipitated,  filter,  and  evaporate  the  filtrate 
until  crystallization  takes  place.  If  during  the  evaporation  any 
more  ferric  hydroxide,  which  may  have  been  held  in  colloidal 
solution,  separates,  filter  it  off  after  the  volume  has  been  reduced 
to  one-half.  Finally,  drain  the  crystals  of  barium  chloride  in  a 
funnel  and  concentrate  the  mother  liquor  for  more  crystals. 
Redissolve  all  the  crystals  in  water,  acidify  with  a  few  drops  of 
hydrochloric  acid,  filter  and  recrystallize.  Yield,  100  to  110  g.  of 
BaCl2-2  H2O. 


68  HALOGEN    COMPOUNDS. 

41.   Manganous  Chloride  from  Waste  Manganese  Liquors. 

Pure  manganous  chloride  can  be  obtained  by  crystallizing  the  waste 
liquor  obtained  in  the  preparation  of  chlorine  from  pyrolusite  and  hydrochloric 
acid,  if,  as  in  the  last  preparation,  the  iron  is  first  removed. 

Similarly,  the  waste  liquor  from  hydrogen  generators  can  be  freed  from 
iron  and  worked  up  into  pure  zinc  salts. 

Evaporate  2  to  3  liters  of  the  manganese  liquor  in  a  porcelain 
dish  over  the  free  flame  in  order  to  expel  the  excess  of  hydro- 
chloric acid.  Dissolve  the  residue,  which  solidifies  on  cooling,  in 
3  to  4  liters  of  water.  Dilute  one-tenth  of  this  solution  in  a  flask 
with  a  large  amount  of  water,  and  add  caustic  soda,  avoiding  an 
excess,  in  order  to  precipitate  manganese  hydroxide.  After 
settling,  siphon  off  the  clear  solution,  shake  up  the  precipitate 
with  pure  water,  and  wash  it  repeatedly  by  decantation  until,  at 
the  end  of  three  or  four  days,  all  of  the  sodium  salt  has  been 
removed.  Add  the  manganese  slime  thus  obtained  to  the  remain- 
ing nine-tenths  of  the  first  solution,  and  allow  the  mixture  to  stand 
for  several  days  in  a  thick-walled,  five-liter  flask  which  is  placed 
in  a  warm  place.  Shake  the  mixture  frequently  until  all  of  the 
iron  has  been  precipitated  by  means  of  the  manganese  hydroxide. 
This  usually  requires  about  two  days,  and  at  the  end  a  little  of 
the  filtered  solution  should  give  no  test  with  potassium  thiocya- 
nate.  Filter  the  solution  through  a  plaited  filter  and  evaporate  it 
until  crystals  begin  to  separate.  Then  transfer  it  to  a  flask  and 
cool  rapidly  while  rotating  under  the  water  tap.  Drain  the  crystal 
meal  in  a  suction-funnel  and  wash  it  first  with  50%,  and  then  with 
pure  alcohol.  By  evaporating  the  mother  liquor  a  further  yield 
is  obtained.  Allow  the  light  pink  crystals,  which  ar,e  still  moist 
with  alcohol,  to  dry  in  contact  with  the  air. 

42.   Anhydrous  Ferric  Chloride;  Preparation  of  Chlorine. 

Clamp  a  tubulated  retort  (cf.  Fig.  15,  p.  78)  of  about  250  c.c. 
capacity  so  that  its  neck  (which  is  1  to  2  centimeters  wide)  is  in  a 
horizontal  position.  Insert  a  bundle  of  iron  wires  (about  0.1  cm. 
in  diameter),  weighing  10  to  25  g.,  to  about  the  middle  of  the  neck, 
and  connect  the  end  of  the  latter,  through  two  sulphuric  acid  wash- 
bottles,  with  a  chlorine  generator.  Into  the  tubulus  of  the  retort 
insert  a  vertical  tube,  about  50  cm.  long  and  1  cm.  wide,  making 


PREPARATION   OF  CHLORINE.  69 

the  joint  tight  by  means  of  a  short  piece  of  rubber  tube.  Place 
the  whole  apparatus  under  the  hood. 

Conduct  a  fairly  rapid  stream  of  chlorine  into  the  retort,  and 
heat  that  part  of  the  neck  which  contains  the  iron  wire  gently  with 
a  small  flame  that  does  not  touch  the  glass.  Very  soon  a  reaction 
begins  to  take  place  with  the  emission  of  light,  and  a  shower  of 
brilliant,  glistening  leaflets  falls  into  the  bulb  of  the  retort.  Heat 
the  neck  of  the  retort  by  fanning  it  with  a  second  and  larger  flame 
until  all  of  the  ferric  chloride  is  sublimed  into  the  bulb.  Tap  the 
vertical  tube  lightly  so  that  any  of  the  product  condensed  in  it 
will  fall  back  into  the  retort. 

At  the  end  of  the  experiment  shake  the  product,  which  is  very 
hydroscopic,  directly  from  the  retort  into  a  dry,  wide-mouthed, 
glass-stoppered  bottle.  The  preparation  keeps  well  if  the  stopper 
is  made  air  tight  with  a  little  vaseline. 

Preparation  of  Chlorine. 

Large  quantities  of  chlorine  may  be  prepared  by  the  oxidation 
of  hydrochloric  acid  with  pyrolusite.  Half  fill  a  round-bottomed 
flask  of  from  1.5  to  2  liters  capacity  (cf.  Fig.  14,  p.  73)  with 
lumps  of  pyrolusite,  and  close  the  flask  with  a  two-holed  stopper; 
through  one  hole  insert  a  thistle  tube,  which  serves  as  a  safety 
tube,  and  must  extend  20  to  30  cm.  above  the  top  of  the  flask 
and  nearly  to  the  bottom  inside.  Through  the  other  hole  insert 
a  short  delivery  tube  bent  at  a  right  angle.  To  cause  the  evolu- 
tion of  chlorine,  pour  concentrated  commercial  hydrochloric  acid 
through  the  thistle  tube  until  the  pyrolusite  is  just  covered,  and 
heat  the  mixture  gently  on  a  Babo  funnel;  regulate  the  rate  at 
which  the  gas  is  generated  by  altering  the  height  of  the  flame 
under  the  evolution  flask.  A  single  charge  of  pyrolusite  is 
sufficient  to  react  with  several  refillings  of  the  acid.  Wash  the 
gas  with  water,  and  dry  it  (if  necessary)  by  passing  it  through 
one  or  two  bottles  containing  concentrated  sulphuric  acid. 

It  is  very  convenient  to  use  the  dry  liquid  chlorine  which 
can  be  purchased  in  steel  cylinders  under  pressure;  but  this 
is  only  to  be  recommended  when  large  amounts  are  frequently 
used. 

When  it  is  desired  to  prepare  a  definite  amount  of  chlorine,  an 
excess  of  hydrochloric  acid  can  be  decomposed  by  means  of  a 


70  HALOGEN    COMPOUNDS. 

weighed   quantity  of  potassium  permanganate  or  of  potassium 
pyrochromate. 

2  KMnO4  +  16  HC1  =  2  MnCl,  +  2  KC1  +  8  H2O  +  5  C12, 
K2Cr207  +  14  HC1  =  2  CrCl3"  +  2  KC1  +  7  H2O  +  3  C12. 

43.   Anhydrous  Ferrous  Chloride ;  Preparation  of  Hydrogen 

Chloride. 

On  account  of  its  higher  melting-point  it  is  more  difficult  to 
prepare  anhydrous  ferrous  chloride  than  the  corresponding  ferric 
salt.  Clamp  a  porcelain  tube,  50  to  60  cm.  long  and  3  cm.  in 
inside  diameter,  in  a  horizontal  position  under  the  hood.  Adjust 
it  at  a  suitable  height  above  the  blast  lamp,  and  surround  it  with 
an  asbestos  heating  chamber  (Fig.  4,  p.  3).  Introduce  a  loose 
bundle  of  iron  wires  (0.1  cm.  in  diameter),  weighing  from 
12  to  15  g.,  into  the  part  of  the  tube  which  can  be  heated 
hottest.  Conduct  into  the  tube  a  rapid  current  of  hydrogen 
chloride  gas  (see  below)  which  is  produced  by  the  action  of 
concentrated,  commercial  sulphuric  acid  upon  750  c.c.  of  con- 
centrated, commercial  hydrochloric  acid.  Place  a  beaker  in  a 
tilted  position  over  the  open  end  of  the  tube.  Maintain  the 
temperature  as  high  as  possible  during  the  experiment.  The 
product  condenses  in  the  cooler  parts  of  the  tube,  and  to  some 
extent  in  the  beaker. 

The  yield  is  15  to  20  g.  of  dirty-white,  hygroscopic,  leaf-like 
crystals. 

Small  amounts  of  ferrous  chloride  can  be  prepared  in  a  similar 
way  in  a  wide  combustion  tube  which  is  heated  in  a  furnace. 

Aluminium  chloride  can  be  prepared  in  like  manner  from 
aluminium  and  hydrogen  chloride. 

Preparation  of  Hydrogen  Chloride. 

Small  amounts  of  hydrogen  chloride  are  most  conveniently 
prepared  in  a  Kipp  generator  containing  large  pieces  of  sal 
ammoniac  (ammonium  chloride)  upon  which  concentrated  sul- 
phuric acid  is  allowed  to  act.  This  method  does  not  work  as 
well  for  preparing  large  quantities  of  the  gas,  because  the  foam- 
ing which  occurs  becomes  troublesome. 


ANHYDROUS   CHROMIUM  TRICHLORIDE.  71 

Larger  amounts  of  hydrogen  chloride  may  be  obtained  by 
treating  200  g.  common  salt  (sodium  chloride)  with  a  cooled 
mixture  of  320  g.  sulphuric  acid  and  80  g.  water  in  a  round- 
bottomed  flask,  which  may  be  conveniently  heated  upon  a  Babo 
boiling  funnel.  On  cooling,  the  flow  of  the  gas  slackens,  but  it 
can  be  started  again  by  renewed  heating.  When  the  mixture  has 
become  exhausted,  it  should  be  poured  out  of  the  flask  while 
still  warm,  since  on  cooling  it  becomes  solid,  and  is  then  difficult 
to  remove. 

The  following  method,  which  depends  upon  the  fact  that  hydro- 
gen chloride  is  but  slightly  soluble  in  concentrated  sulphuric 
acid,  is  also  to  be  recommended.1  Allow  concentrated  sulphuric 
acid  to  drop  from  a  dropping  funnel  into  a  large  suction  flask 
containing  concentrated  hydrochloric  acid,  to  which  a  handful  of 
common  salt  may  with  advantage  be  added.  It  is  advisable  to 
let  the  sulphuric  acid  fall  first  into  a  small  test-tube  so  that  by 
its  flowing  uniformly  over  the  edge  of  the  latter  a  steady  evolu- 
tion of  gas  is  produced.  The  rate  of  flow  of  gas  from  the  flask 
is  governed  by  the  stopcock  which  regulates  the  dropping  of  the 
sulphuric  acid. 

The  reverse  process  of  adding  hydrochloric  acid  to  concen- 
trated sulphuric  acid  may  also  be  used,  but  in  order  to  make  the 
lighter  hydrochloric  acid  solution  mix  thoroughly  with  the  heavier 
sulphuric  acid  it  is  necessary  to  draw  out  the  stem  of  the  dropping 
funnel  into  a  capillary  which  reaches  nearly  to  the  bottom  of  the 
generating  flask. 

44.   Anhydrous  Chromium  Trichloride. 

Place  10  to  20  g.  of  coarsely  powdered,  metallic  chromium 
(No.  3)  in  an  apparatus  constructed  like  that  used  in  the  last 
preparation.  Expel  the  air  completely  by  means  of  a  stream  of 
dry  chlorine,  which  will  take  at  least  half  an  hour,  and  then  heat 
the  metal  in  a  current  of  perfectly  dry  chlorine  for  from  thirty 
minutes  to  an  hour  at  as  high  a  temperature  as  possible.  After 
cooling,  replace  the  chlorine  by  carbon  dioxide  and  shake  the 
preparation  out  of  the  tube.  Beautiful,  glistening,  violet  flakes 

1  The  same  method  may  be  used  for  the  production  of  sulphur  dioxide  if 
concentrated  sulphuric  acid  is  allowed  to  drop  into  a  solution  of  commercial 
sodium  bisulphite. 


72  HALOGEN  COMPOUNDS. 

are  obtained,  together  with  a  darker  powder  of  a  more  brownish 
shade  which  under  the  microscope  is  shown  to  be  crystalline. 
Collect  the  two  forms  separately.  Chromium  trichloride  is,  on 
account  of  its  slow  rate  of  solution,  practically  insoluble  in  water. 
It  is  difficult  to  clean  the  porcelain  tube  after  this  experiment. 
Stopper  it  at  one  end,  and  fill  it  with  commercial,  concentrated 
hydrochloric  acid  to  which  a  little  ferrous  sulphate  has  been 
added,  and  let  it  stand.  Dependent  preparations:  Chromium 
Nitride  (No.  62),  Hexamminechromic  Nitrate  and  Chloropen- 
tamminechromic  Chloride  (No.  140). 

45.  Sulphur  Chloride,  S2C12. 

Of  the  chlorides  of  sulphur,  S2C12  is  stable  at  the  laboratory  temperature. 
The  chloride  SC14  is  a  yellowish-white  substance  which  melts  at  —  30°  and 
decomposes  at  a  few  degrees  above  the  melting-point.  The  compound  SC12 
has  very  recently  been  prepared  pure  by  E.  Beckmann  in  the  form  of  a  dark- 
red  substance,  and  has  been  shown  by  him  to  be  a  definite  compound. 

Connect  a  tubulated  retort  (Fig.  14)  of  250  c.c.  capacity  with 
a  condenser,  an  adapter,  and  a  suction  flask  which  serves  as  a 
receiver;  from  the  side  arm  of  the  latter,  lead  a  glass  tube  to 
the  ventilating  flue.  Fill  the  retort  with  100  g.  of  flowers  of 
sulphur,  and  heat  it  upon  a  Babo  funnel,  or,  better  still,  in  a 
nickel  air-bath,  to  a  temperature  of  from  200°  to  250°.  Conduct 
into  the  retort  a  rapid  current  of  chlorine,  which  has  been 
washed  once  with  water  and  once  with  sulphuric  acid.  The  sul- 
phur chloride  that  is  formed  distils  completely  into  the  receiver; 
a  slight  blackish  coating  which  remains  behind  is  due  to  impurities 
in  the  sulphur.  The  operation  requires  about  four  hours. 

To  the  crude  product  thus  prepared,  add  10  to  15  g.  of  sulphur 
which  combines  with-  the  excess  of  chlorine.  Distil  the  sulphur 
chloride  from  a  fractionating  flask  provided  with  a  condenser 
and  a  receiver,  and  if  more  than  a  small  amount  passes  over 
before  a  temperature  of  130°  is  reached,  pour  it  back  into  the 
flask  and  add  more  sulphur.  Then  distil  again,  rejecting  the 
portion  that  passes  over  below  134°,  and  collecting  as  pure  pro- 
duct the  distillate  between  137°  and  138°.  Only  towards  the 
end  of  the  process  does  the  temperature  rise  a  few  degrees  above 
this  point,  due  to  superheating  of  the  vapors.  The  excess  of 
sulphur  remains  behind  in  the  flask.  The  preparation  may  be 


CHLORIDES   OF  PHOSPHORUS. 


73 


further  purified  by  another  fractionation.     The  yield  is  almost 
theoretical. 

Sulphur  chloride,  S2C12,  as  ordinarily  obtained,  is  a  yellowish- 
red,  heavy  liquid  which  fumes  a  little  in  the  air,  has  an  unpleas- 
ant odor,  and  attacks'  the  mucous  membrane.  When  distilled  in 
vacuum  it  has  a  pure  yellow  color.  This  preparation  should  be 


Fig.  14 

carried  out  in  a  room  reserved  especially  for  working  with  nox- 
ious compounds,  and  the  liquids  should  be  transferred  from  one 
vessel  to  another  only  when  under  the  hood. 

46.    Chlorides  of  Phosphorus. 

Chlorine  combines  with  phosphorus  to  form  either  the  liquid  trichloride  or 
the  solid  pentachloride  according  to  whether  the  phosphorus  or  chlorine  is 
present  in  excess.  Phosphorus  trichloride  boils  undecomposed  at  76°  C. ;  on 
heating  phosphorus  pentachloride,  it  vaporizes  without  passing  through  the 
liquid  phase.  Density  determinations  of  the  gas  show  that  the  pentachloride 
is  appreciably  dissociated  at  180°,  and  practically  completely  so  at  temper- 
atures above  290°,  into  phosphorus  trichloride  and  chlorine;  the  latter  can  be 
identified  by  its  yellowish-green  color.  On  cooling,  the  dissociation  products 
recombine  and  again  form  the  pentachloride. 

Phosphorus  oxychloride  is  most  conveniently  prepared  by  oxidizing  phos- 
phorus trichloride  with  potassium  chlorate. 

3  PC13  +  KC103  =  3  POC13  +  KC1. 


74  HALOGEN  COMPOUNDS. 

Phosphorus  Trichloride.  Use  the  same  apparatus  as  in  the 
preparation  of  sulphur  chloride  (No.  45),  except  that  the  heating 
bath  under  the  retort  is  in  this  case  superfluous.  Cut  31  g.  of 
yellow  phosphorus  into  small  pieces,  keeping  it  under  water  in  a 
porcelain  mortar,  and  handling  it  with  pincers.  Dry  the  phos- 
phorus by  rapidly  pressing  each  piece  between  filter  papers,  and 
introduce  it  through  the  tubulus  into  the  retort  which  has  pre- 
viously been  filled  with  carbon  dioxide  gas.  Avoid  handling  the 
phosphorus  with  the  fingers,  because  it  takes  fire  easily,  and 
phosphorus  burns  are  severe  and  frequently  dangerous.  Any 
residue  of  phosphorus  left  in  the  mortar  should  be  wiped  out  with 
moist  filter  paper,  and  the  paper  immediately  burned. 

On  conducting  chlorine  into  the  retort,  the  phosphorus  ignites, 
melts,  and  burns  with  a  pale  flame  to  phosphorus  trichloride. 
Outside  heating  of  the  retort  is  unnecessary,  but  it  is  important 
to  have  the  stream  of  chlorine  pass  rapidly  and  steadily.  Regu- 
late the  reaction  from  time  to  time,  as  necessary,  by  raising  or 
lowering  the  tube  through  which  the  chlorine  enters  the  retort; 
this  tube  should  be  fitted  so  that  it  can  be  moved  readily. 
Lower  the  tube  if  a  white  sublimate  of  phosphorus  pentachloride 
forms  in  the  upper  part  of  the  neck  of  the  retort;  raise  it  a 
little  if  a  yellowish-red  sublimate  begins  to  make  the  neck  of 
the  retort  opaque.  This  regulation,  which  is  necessary  for  the 
proper  carrying  out  of  the  experiment,  does  not  involve  any 
difficulty. 

Purify  the  crude  product  by  distilling  it  from  a  flask  with  a 
side-arm  condenser  (Fig.  7,  p.  6).  Boiling-point,  76°.  Yield,  100 
to  120  g. 

Phosphorus  Pentachloride.  Close  a  wide-mouthed  liter  bottle 
by  means  of  a  three-holed  cork.  Through  one  hole  introduce  the 
stem  of  a  dropping  funnel  so  that  it  reaches  just  inside  the 
stopper;  through  the  second  hole  insert  a  glass  tube  of  1  cm. 
bore  which  reaches  to  the  middle  of  the  bottle  and  serves  for  the 
introduction  of  chlorine;  through  the  third  hole  insert  another 
tube  which  ends  just  below  the  cork  and  serves  for  the  escape  of 
the  excess  of  chlorine.  Provide  the  1  cm.  wide  vertical  tube  with  a 
side  arm  through  which  the  chlorine  is  to  be  introduced,  and  close 
the  upper  end  with  a  cork  stopper.  Then  when  the  lower  part  of 
the  tube  becomes  clogged  with  phosphorus  pentachloride,  it  is 


CHLORIDES   OF   ANTIMONY.  75 

only  necessary  to  remove  the  stopper  a  moment  and  push  the 
obstruction  out  of  the  way  with  a  stirring  rod. 

Conduct  chlorine  into  the  flask,  and  at  the  same  time  intro- 
duce phosphorus  trichloride  through  the  dropping  funnel;  the 
two  substances  immediately  combine.  Care  should  be  taken  to 
keep  chlorine  present  in  excess.  At  the  end  loosen  the  phos- 
phorus pentachloride  formed  by  means  of  a  spatula,  and  allow 
the  flask,  which  is  still  filled  with  chlorine,  to  stand  for  some  time 
before  removing  the  product.  The  yield  is  almost  quantitative. 
Dependent  preparation:  Thioayl  Chloride,  No.  151. 

Phosphorus  Oxy chloride.  Connect  a  150  to  200  c.c.  distilling 
flask,  containing  21  g.  (one-sixth  mol.)  of  finely  powdered  potas- 
sium chlorate,  with  a  condenser  and  receiver.  Then  allow  69  g. 
of  phosphorus  trichloride  (one-half  mol.)  to  flow  into  the  flask  a 
little  at  a  time.  After  each  addition  of  the  chloride,  wait  until 
the  reaction,  which  is  made  evident  by  a  gentle  ebullition,  has 
ceased  before  adding  more;  at  the  start  it  is  permissible  to 
warm  slightly  if  necessary.  Should  a  little  liquid  distil  over  into 
the  receiver  during  this  operation,  return  it  to  the  distilling 
flask. 

When  the  reaction  is  complete,  distil  the  phosphorus  oxychlo- 
ride  by  heating  with  a  large  flame,  holding  the  burner  in  the  hand 
and  playing  the  flame  around  the  bulb  of  the  flask.  A  ther- 
mometer is  not  necessary  for  this  distillation.  Clean  and  dry 
the  apparatus,  and  redistil  the  product,  this  time  using  a  ther- 
mometer. Collect  the  first  few  drops  which  come  over  separately. 
Boiling-point,  110°.  Dependent  preparation:  Triethyl  Phos- 
phate, No.  157. 

47.    Chlorides  of  Antimony. 

Antimony  Trichloride.  Treat  100  g.  of  finely  powdered  stibnite, 
in  a  750  c.c.  flask,  with  400  g.  of  concentrated,  commercial  hydro- 
chloric acid,  shaking  frequently  and  heating  upon  the  water  bath 
until  as  much  as  possible  of  the  material  is  dissolved.  Boil  the 
solution  five  minutes  in  order  to  remove  the  greater  part  of  the 
dissolved  hydrogen  sulphide.  Add  five  cubic  centimeters  more  of 
concentrated  hydrochloric  acid,  and  filter  the  solution  through  a 
Biichner  funnel  containing  a  layer  of  asbestos-felt  which  has  been 
previously  moistened  with  the  concentrated  acid. 


76  HALOGEN    COMPOUNDS. 

Distil  the  antimony  trichloride  solution  from  a  retort,  provided 
with  a  thermometer,  until  the  temperature  reaches  120°;  in  order 
to  prevent  bumping,  place  bits  of  pumice  or  of  unglazed  porcelain 
in  the  liquid.  The  distillate  consists  chiefly  of  hydrochloric  acid, 
containing  eventually  some  arsenic  trichloride. 

Transfer  the  liquid  remaining  in  the  retort,  after  again  filtering 
through  asbestos  if  necessary,  to  a  distilling  flask.  Fit  the  side 
arm  of  the  latter  by  means  of  a  cork  stopper  to  a  long  tube 
about  1  cm.  in  diameter  which  serves  as  a  condenser  (cf.  Fig.  8, 
p.  7).  The  distillate  first  passing  over  is  clear,  then  becomes 
more  or  less  yellowish  due  to  ferric  chloride,  and  finally,  when  the 
temperature  is  above  215°,  it  becomes  colorless  again.  At  this 
point  change  the  receiver  for  a  dry,  clean,  weighed  Erlenmeyer 
flask.  The  last  portion  of  the  distillate  solidifies,  on  cooling,  into 
a  radiating  mass  of  crystals.  It  can  be  further  purified  by  redis- 
tillation. Boiling-point,  223°. 

Pour  the  first  fraction  of  the  distillate,  which  consists  of  a  mix- 
ture of  hydrochloric  acid  and  antimony  trichloride,  into  a  large 
quantity  of  water;  collect  the  precipitated  basic  antimony  chloride 
(essentially  a  mixture  of  SbOCl  and  Sb2O3)  on  a  filter  and  wash 
and  dry  it.  Dependent  preparation:  Metallic  Antimony,  No.  7. 

Antimony  Pentachloride.  Pass  dry  chlorine  gas  into  fused 
antimony  trichloride  until  the  gain  in  weight  corresponds  to  that 
required  for  the  change  to  the  pentachloride.  At  first  the  reaction 
mixture  must  be  kept  above  the  melting-point  of  pure  trichloride, 
but  as  more  and  more  of  the  pentachloride,  which  is  liquid  under 
ordinary  conditions,  forms,  the  mixture  may  be  allowed  to  cool  to 
room  temperature. 

48.   Iodides  of  Bismuth. 

Bismuth  Tri-iodide.  Triturate  8  g.  of  sifted  powdered  bismuth 
with  13  g.  of  iodine  in  a  mortar,  and  introduce  the  mixture  into  a 
50  c.c.  plain  retort.  Cut  off  the  neck  of  the  retort  to  a  length  of 
about  7  cm.,  and  suspend  the  whole  with  a  loop  or  spiral  of  wire 
from  a  ring-stand  in  such  a  way  that  all  portions  of  the  retort  can 
be  heated  freely.  On  heating  the  mixture  a  feeble  reaction  is  soon 
observed.  On  heating  more  strongly,  a  little  iodine  sublimes  at 
first,  and  should  be  driven  off  by  playing  a  second  flame  over  the 
neck  of  the  retort  (Hood);  then  the  bismuth  tri-iodide  sublimes 


BISMUTH  TRIBROMIDE.  77 

and  condenses  in  the  form  of  a  shower  of  crystalline  spangles. 
Collect  these  in  a  porcelain  evaporating  dish  which  is  placed  so 
that  the  neck  of  the  retort  rests  in  its  lip.  Cover  the  dish  with  a 
watch-glass,  and  protect  the  space  still  left  open  with  some  asbestos 
paper  cut  to  the  proper  shape. 

On  standing,  bismuth  iodide  decomposes  rather  easily  with  lib- 
eration of  iodine. 

Basic  Bismuth  Iodide,  BiOL  Triturate  10  g.  of  bismuth  tri- 
iodide  in  a  mortar  with  water,  and  decant  off  the  liquid  together 
with  the  finest  powder  into  a  beaker;  treat  the  residue  with 
another  portion  of  water  in  exactly  the  same  way,  and  continue 
the  treatment  until  all  of  the  material  is  obtained  in  a  state  of 
finest  subdivision  with  from  200  to  400  c.c.  of  water.  Boil  the 
mixture  for  an  hour  or  two,  collect  the  product  of  hydrolysis  on  a 
suction  filter,  wash  it  with  water,  and  dry  it  in  the  hot  closet. 
Small,  light-brown  to  red  crystal  leaflets  are  obtained  which  are 
somewhat  lighter  colored  than  red  phosphorus.  Yield,  about  7.5  g. 

The  hydrolysis  of  the  iodide,  which  is  difficultly  soluble  in  water,  takes 
place  more  slowly  than  that  of  the  more  soluble  chloride,  an  indication  that 
such  interactions  of  a  solid  substance  under  a  liquid  do  not  take  place  with 
the  solid  itself,  but  involve  rather  only  that  part  of  the  substance  which  exists, 
at  the  moment,  in  solution. 

49.   Bismuth  Tribromide. 

A  mixture  of  powdered  bismuth  and  bromine  reacts  only  slowly  and  in- 
completely in  the  cold,  but  with  almost  explosive  violence  when  heated.  The 
simplest  way  to  prepare  bismuth  tribromide  is  to  pass  bromine  vapors  over 
heated  bismuth. 

The  apparatus  employed  is  similar  to  that  used  in  the  prepara- 
tion of  ferric  chloride  (No.  42).  The  neck  of  a  tubulated  retort, 
of  from  100  to  150  c.c.  capacity,  is  bent  downward  a  little,  near 
the  bulb,  and  a  little  farther  away  it  is  bent  upward  (cf.  Fig.  15). 
Introduce  10  to  15  g.  of  bismuth  in  the  depression  of  the  neck, 
and  connect  the  end  of  the  latter  with  a  wash  bottle  containing 
bromine,  through  which  dry  carbon  dioxide  is  to  be  passed  in  the 
direction  of .  the  retort.1  Through  the  tubulus  of  the  retort 

1  Instead  of  conducting  bromine  diluted  with  carbon  dioxide  over  the 
bismuth,  the  pure  vapors  may  be  distilled  from  a  second  smaller  retort.  The 
reaction  then  takes  place  more  rapidly,  but  requires  more  attention. 


78 


HALOGEN  COMPOUNDS. 


introduce  a  glass  tube  one-half  meter  long  and  about  1  cm.  in 
diameter,  and  make  the  joint  tight  with  a  piece  of  rubber  tubing 
around  the  glass.  This  long  tube  is  to  serve  as  a  condenser,  and 
to  lead  the  excess  of  bromine  into  the  hood. 

After  starting  the  current  of  carbon  dioxide,  heat  the  bismuth 
gently  by  means  of  a  Bunsen  burner  with  a  flame  spreader,  and 
warm  the  bromine  by  placing  the  wash  bottle  in  warm  water  and 


FIG.  15. 


renewing  the  latter  as  it  cools.  Dark-red  vapors  of  bismuth 
bromide  form  and  condense  as  yellow  flakes  in  the  bulb  of  the 
retort,  and  to  some  extent  in  the  glass  tube. 

Break  the  retort,  collect  the  product,  and  let  it  stand  over  night 
in  a  vacuum  desiccator  to  remove  any  adhering  bromine ;  or,  distil 
it  from  a  smaller  plain  retort.  The  neck  of  the  latter  should  be- 
about  8  cm.  long  and  the  whole  should  be  suspended  in  a  double 
loop  of  wire  (cf.  No.  48).  Yield,  nearly  theoretical. 

Boiling-point  of  Bismuth  Bromide.  Use  a  thermometer  filled 
under  pressure,  which  measures  temperatures  as  high  as  540°. l 
Place  the  bismuth  bromide  in  a  Jena  glass  test-tube,  20  cm.  long 
and  2.5  cm.  wide,  and  heat  until  the  substance  boils;  suspend 
the  thermometer  so  that  it  reaches  well  into  the  vapor.  When  the 
vapors  first  begin  to  reach  the  mercury  bulb,  lower  or  remove  the 

1  Cf .  No.  54  for  the  thermoelectric  measurement  of  temperatures. 


TIN  TETRACHLORIDE  79 

flame  for  a  little  so  as  to  avoid  heating  the  thermometer  too 
suddenly;  then  heat  strongly  again.  Correct  the  boiling-point  as 
thus  determined  either  as  directed  in  No.  6,  or  by  means  of  finding 
the  apparent  boiling-point  of  pure  sulphur  under  exactly  the  same 
conditions.  The  difference  between  the  latter  observed  reading 
and  448°  (the  true  boiling-point  of  sulphur)  is  to  be  applied  as  a 
correction  to  the  apparent  boiling-point  of  the  bismuth  tribromide. 
This  very  simple  method  of  making  the  correction  for  stem  expo- 
sure gives  good  results,  provided  the  two  temperatures  compared 
are  not  far  from  one  another.  Boiling-point  of  bismuth  bromide, 
4660.1 

50.   Tin  Tetrachloride.2 

Make  use  of  the  same  apparatus  as  described  under  the  prep- 
aration of  sulphur  chloride  (Fig.  14,  p.  73).  Place  60  g.  of  tin 
in  the  retort,  and  heat  it  by  placing  a  Bunsen  flame  underneath. 
After  the  tin  is  melted  pass  a  rapid  current  of  chlorine,  allowing 
it  to  play  directly  on  the  surface  of  the  metal.  Liquid  tin 
tetrachloride,  which  is  colored  yellow  by  dissolved  chlorine^ 
collects  in  the  receiver.  When  the  metal  in  the  retort  has 
all  disappeared,  add  some  tinfoil  to  the  distillate  in  the  receiv- 
ing flask,  stopper  the  latter,  and  allow  it  to  stand  until  the 
next  day  in  order  that  the  dissolved  chlorine  may  all  react  with 
the  tin. 

Place  the  crude  product  together  with  a  little  tinfoil  in  a  100  c.c. 
distilling  flask,  provided  with  a  side-arm  condenser  (Fig.  7).  Fit 
a  thermometer  in  the  neck  of  the  flask,  and  distil  the  liquid  at 
the  hood,  rejecting  the  first  few  drops  which  pass  over.  If  the 
distillate  is  not  perfectly  colorless,  allow  it  to  stand  over  night  with 
more  tinfoil  and  repeat  the  distillation.  Boiling-point,  113.5°  to 
114°.  Preserve  the  preparation  in  a  sealed  vessel. 

When  exposed  to  the  atmosphere,  tin  tetrachloride  absorbs 
water  and  soon  becomes  changed  to  the  solid  white  hydrate. 
Tin  tetrachloride  can,  like  titanium  tetrachloride,  be  transformed 


1  Author's  observation:  Victor  Meyer  found  453°  with  an  air  thermometer. 
Ann.  264,  122  (1891). 

2  For  a  somewhat  different  method  of  carrying  out  this  preparation,  see 
R.  Lorenz,  Z.  anorg.  Chem.,  10,  44  (1895). 


80  HALOGEN   COMPOUNDS. 

into  the  corresponding  sulphide  (cf.  Titanium  Bisulphide,  No.  58). 
The  formation  of  colloidal  stannic  acid  from  tin  tetrachloride 
is  discussed  under  Purple  of  Cassius,  No.  25. 

51.    Silicon  Tetrachloride. 

For  preparing  silicon  tetrachloride  from  the  elements,  either  commercial 
silicon  obtained  in  the  electric  furnace,  or  the  crystallized  product  prepared 
by  the  thermite  process  (No.  4),  may  be  used. 

Generate  chlorine  in  a  two-liter  (or  larger)  flask  from  pyrolu- 
site  and  concentrated  hydrochloric  acid,  and  wash  the  gas  once 
with  water  and  twice  with  concentrated  sulphuric  acid.  Spread 
a  layer  of  about  10  g.  finely  powdered  silicon  loosely  in  a  40  cm. 
long  combustion  tube  which  is  placed  over  a  row  burner  (Fig.  3). 
Connect  one  end  of  the  combustion  tube  with  the  chlorine  gener- 
ator, and  draw  out  the  other  end  to  about  the  size  of  a  lead 
pencil.  Join  this  narrow  end  with  a  gas  wash-bottle,  using  a 
rubber  connector  and  pushing  the  ends  of  the  glass  tubes  close 
together.  Cool  the  wash  bottle  by  surrounding  it  with  a  mix- 
ture of  ice  and  salt,  and  arrange  a  glass  tube  to  conduct  the 
waste  gases  into  the  ventilating  flue. 

First  of  all,  —  and  this  is  very  important, — sweep  the  air  com- 
pletely out  from  the  apparatus  by  passing  a  rapid  stream  of 
chlorine  gas  for  about  half  an  hour.  After  that,  heat  the  com- 
bustion tube  until  the  reaction  begins  and  produces  incandes- 
cence; the  flames  beneath  the  tube  may  be  turned  quite  low 
while  the  reaction  is  progressing,  and  it  is  well  to  turn  the  tube 
from  time  to  time  on  its  long  axis.  All  of  the  silicon  is  acted 
upon,  and  only  a  few  flakes  of  silicon  dioxide  remain  behind, 
while  a  trace  of  aluminium  chloride  condenses  at  the  end  of  the 
tube. 

Without  using  a  thermometer,  distil  the  impure  product  slowly 
from  a  fractionating  flask  with  side-arm  condenser  (Fig.  7). 
The  greater  part  of  the  dissolved  chlorine  is  thereby  expelled, 
but  to  remove  the  last  of  it,  let  the  distillate  stand  about  a  day 
in  contact  with  mercury  in  a  thick-walled  bottle,  stoppered  with 
a  cork  (not  a  glass  stopper),  shaking  vigorously  from  time  to  time 
until  the  liquid  is  decolorized.  By  again  distilling,  this  time 
with  a  thermometer,  the  compound  is  obtained  pure.  Boiling- 
point,  58°  to  60°.  Yield,  35  to  40  g. 


TITANIUM  TETRACHLORIDE.  81 

That  portion  of  the  distillate  passing  over  above  60°,  which, 
however,  is  small  in  amount  by  this  method  of  preparation,  con- 
tains a  little  silicon  hexachloride,  Si2Cle,  boiling-point  145°  to 
146°,  and  some  silicon  octachloride,  Si3Cl8,  boiling-point  210° 
to  215°. 

Silicon  tetrachloride  is  a  colorless,  mobile  liquid,  and  shows 
a  high  refractive  index  for  light.  It  fumes  strongly  in  the  air, 
and  on  being  mixed  with  water  it  hydrolyzes,  forming  ortho- 
silicic  and  hydrochloric  acids: 

SiCl4  +  4  H2O  =  Si(OH)4  +  4  HC1. 

Preserve  the  preparation  in  a  sealed  flask.  Dependent  prepara- 
tion: Tetraethyl  Silicate,  No.  158. 

Instead  of  starting  with  pure  silicon,  the  directions  of  Gatter- 
mann  *  may  be  followed :  Prepare  an  impure  silicon  by  igniting 
40  g.  magnesium  powder  with  160  g.  dried  and  sifted  quartz 
sand,  and  chlorinate  the  resulting  mixture  of  silicon  and  mag- 
nesium oxide  by  heating  it  to  300°  to  310°  in  a  long  combustion 
tube,  placed  in  a  "  bomb  "  furnace,  while  passing  a  current  of  dry 
chlorine.  This  method  yields  a  product  containing  more  of  the 
hexachloride  and  octachloride  than  when,  as  by  the  first  method, 
the  chlorination  is  carried  out  at  a  higher  temperature. 

Silicon  Chloroform.  The  compound  SiHCl3  can  be  obtained  in 
a  corresponding  manner  if,  instead  of  chlorine,  dry  hydrogen 
chloride,  free  from  air,  is  passed  over  the  silicon  powder  at  a 
temperature  of  450°  to  500°.  For  the  preparation  of  this  com- 
pound in  larger  quantities,  and  for  its  properties  (boiling  point, 
33°  to  34°),  see  Ruff  and  Albert,  Ber.  38,  2222  (1905). 

52.   Titanium  Tetrachloride  from  Rutile. 

A  number  of  oxides  which  are  not  reduced  by  charcoal  can  be  transformed 
into  the  corresponding  chlorides  by  the  simultaneous  action  of  carbon  and 
chlorine.  This  method  was  originated  by  Oersted  in  1824,  and  perfected  by 
Wohler  for  a  number  of  different  oxides.  It  has  been,  for  decades,  the  most 
important,  if  not  the  sole,  method  for  the  preparation  of  certain  chlorides 
(A1C13,  SiCl4,  TiCl4,  UC14  (No.  169),  etc.).  From  a  theoretical  standpoint  it  is 
a  good  example  of  the  displacement  of  an  equilibrium  by  the  removal  of  one 
or  more  of  the  products  from  the  sphere  of  action,  whereby  a  given  reaction  is 
enabled  to  become  quantitative:  the  reduction  of  aluminium  oxide  cannot  be 

1  L.  Gattermann,  Ber.  22,  186;  (1889)  27,  1943  (1894). 


82  HALOGEN    COMPOUNDS. 

accomplished  by  means  of  charcoal,  but  the  oxides  of  carbon,  on  the  other 
hand,  can  be  reduced  by  aluminium;  an  equilibrium  mixture  of  aluminium 
oxide,  aluminium,  carbon,  and  oxides  of  carbon,  therefore,  can  contain  only 
an  infinitesimal  amount  of  aluminium.  If,  however,  this  minute  amount  of 
aluminium  in  the  mixture  is  removed  continuously  by  causing  it  to  combine 
with  chlorine,  an  opportunity  is  thus  afforded  for  a  continuous  reproduction 
of  fresh  quantities  of  aluminium  from  the  reduction  of  its  oxide,  until  the  entire 
amount  of  the  latter  is  exhausted. 

The  preparation  of  titanium  tetrachloride  from  rutile  is  the  most  con- 
venient laboratory  method  for  obtaining  pure  titanium  compounds  from  that 
inexpensive  mineral.  Instead  of  using  carbon  and  chlorine  separately,  vapors 
of  carbon  tetrachloride  may  be  conducted  over  the  heated  oxide;  the  same 
end  may  be  attained  by  using  sulphur  chloride. 

Mix  100  g.  of  finely  powdered  rutile  intimately  with  40  g.  of 
lampblack,  and  knead  the  mixture,  with  the  aid  of  as  little  starch 
paste  as  possible,  into  a  thick,  although  still  plastic  mass.  Shape 
the  mass  into  pellets  of  about  0.5  cm.  diameter,  and  dry  them 
first  in  the  hot  closet,  and  then  in  a  crucible  placed  in  the 
charcoal  furnace.  Introduce  the  dry  and  very  brittle  pellets 
carefully  into  a  wide  combustion  tube,  and  arrange  the  appa- 
ratus, which  must  be  carefully  dried,  as  shown  in  the  sketch 
below  (Fig.  16).  Rubber  connectors  are  to  be  avoided  as  much 


Combustion  Furnace 


FIG.  16. 

as  possible.  Before  beginning  the  chlorination  ignite  the  pellets 
again,  this  time  in  an  atmosphere  of  carbon  dioxide,  in  order  to 
remove  the  last  traces  of  moisture;  and  meanwhile  start  the  evo- 
lution of  chlorine  in  the  generator.  Then  connect  the  apparatus 
for  the  first  time  with  the  receiver  (capacity  200  to  300  c.c.), 
which  is  surrounded  by  ice.  Conduct  chlorine  into  the  appa- 
ratus, and  heat  the  tube,  which  rests  in  a  combustion  furnace,  at 
first  with  small  flames,  then  gradually  bring  it  to  a  red  heat.  It 


TITANIUM  TETRACHLORIDE.  83 

is  essential  to  maintain  a  strong,  steady  stream  of  chlorine,  the 
bubbles  of  which  pass  through  the  wash  bottle  so  rapidly  that  it 
is  just  impossible  to  count  them.  Titanium  tetrachloride  collects 
in  the  receiver  in  the  form  of  a  liquid  which  is  colored  yellow  by 
free  chlorine  and  is  clouded  by  small  crystals  of  ferric  chloride.1 
The  operation  requires  about  three  hours.  The  amount  of  rutile 
prescribed  is  sufficient  to  yield  enough  material  for  two  fillings  of 
a  combustion  tube. 

Filter  the  crude  product  through  a  dry  Gooch  crucible  contain- 
ing an  asbestos-felt,  and  then  allow  it  to  remain  in  contact  with 
mercury,  or  copper  filings,  in  a  thick-walled  bottle  closed  with  a 
cork  stopper;  occasionally  shake  it  vigorously,  adding  a  little  fresh 
metal  each  time,  until,  at  the  end  of  about  twelve  hours,  all  the 
free  chlorine  has  combined.  A  little  sodium  amalgam  may  be 
added  toward  the  end  of  the  process  in  order  to  remove  traces 
of  vanadium.  After  again  filtering  through  a  Gooch  crucible, 
distil  the  clear  liquid  from  a  fractionating  flask  with  a  side-arm 
condenser  (Fig.  7,  p.  6).  Boiling-point,  136°  to  137°.  The  liquid 
has  a  strong  refracting  power  for  light,  and  it  fumes  in  the  air. 
Yield,  100  to  130  g.  Dependent  preparations:  Titanium  Bisul- 
phide, No.  58;  Potassium  Titanium  Fluoride,  No.  103. 

Titanium  Dioxide  from  Titanium  Tetrachloride.  Pour  about 
10  c.c.  of  titanium  tetrachloride,  a  little  at  a  time,  into  200  c.c.  of 
water,  whereby  hydrolysis  takes  place,  although  the  hydrated 
titanium  oxide  which  is  formed  remains  in  colloidal  solution. 
Add  a  little  sulphuric  acid,  and  boil;  titanic  acid  is  thereby 
precipitated  in  a  form  which  can  be  readily  filtered.  Wash  the 
precipitate  with  water  containing  ammonium  nitrate,  dry  it  in 
the  hot  closet,  and  ignite  it  in  a  porcelain  crucible  until  white 
titanium  dioxide  is  left.  In  order  to  remove  the  last  traces  of 
sulphuric  acid,  which  are  retained  very  persistently,  ignite  the 
product  several  times  with  ammonium  carbonate.  Yield,  almost 
theoretical. 

Hydrogen  Peroxide  Reaction.  Fuse  a  little  titanium  dioxide 
with  potassium  acid  sulphate,  and  dissolve  the  melt  in  cold  water. 
A  few  drops  of  this  solution  when  treated  with  an  aqueous  solution 
of  hydrogen  peroxide  show  a  brownish-yellow  to  yellow  color. 

1  If  the  rutile  used  has  a  high  iron  content,  it  can  happen  that  the  tube 
will  become  clogged  with  crystals  of  ferric  chloride. 


84  HALOGEN    COMPOUNDS. 

The  reaction  is  extremely  sensitive,  and  can  be  used  as  a  test  for 
hydrogen  peroxide  as  well  as  for  titanium.  A  drop  of  ordinary  3% 
hydrogen  peroxide  solution  diluted  with  half  a  liter  of  water  can 
be  detected  by  adding  a  few  cubic  centimeters  of  the  reagent. 
Molybdenum  and  vanadium  compounds  react  similarly,  though 
less  strongly,  with  hydrogen  peroxide. 

53.   Anhydrous  Titanium  Trichloride. 

Titanium  tetrachloride  in  aqueous  solution  may  be  reduced  by  metallic 
tin  or  zinc,  or  by  electrolysis,  to  violet  titanium  trichloride;  solutions  of  the 
latter  have  recently  been  recommended  highly  as  reducing  agents.  The 
trichloride  is  obtained  in  an  anhydrous  condition  by  a  method  which  is  of 
quite  general  applicability.  A  mixture  of  titanium  tetrachloride  vapor  and 
hydrogen  is  conducted  through  a  red-hot  tube  whereby  the  trichloride  deposits 
in  the  form  of  reddish-violet,  non-volatile  leaflets. 

Place  a  tube  made  of  difficultly  fusible  glass  in  a  combustion 
furnace,  and  connect  it,  as  shown  in  Fig.  17,  on  the  one  side  with  a 


Combustion  Furnace 


FIG.  17. 

100  c.c.  retort,  and  on  the  other  side  with  a  receiver  of  twice  that 
size.  Rest  the  receiver  in  a  dish  filled  with  ice,  and  add  40  to  50  g. 
of  titanium  tetrachloride  to  the  retort. 

Fill  the  entire  apparatus  with  hydrogen  which  is  conducted 
through  the  inlet  tube  of  the  retort.  As  soon  as  the  gas  escaping 
at  the  other  end  is  shown,  on  testing,  to  consist  of  pure  hydrogen, 
heat  the  combustion  tube  to  bright  redness  and  heat  the  tita- 
nium tetrachloride  nearly  to  the  boiling-point  by  means  of  a  small 
flame,  meanwhile  passing  a  rapid  stream  of  hydrogen  continuously 
through  the  apparatus.  Regulate  the  temperature  of  the  titanium 
tetrachloride  so  that,  as  nearly  as  possible,  all  of  it  is  decomposed 
and  none  condenses  unchanged  in  the  receiver.  If,  however,  this 
is  not  accomplished,  pour  back  the  distillate  into  the  retort  and 
repeat  the  process.  Finally,  disconnect  the  receiver  (close  the 
end  of  the  tube  with  a  cork  carrying  a  short  delivery  tube  bent 
downward)  and  heat  the  part  of  the  tube  projecting  beyond  the 
furnace  by  fanning  it  with  a  flame,  until  no  more  white  vapors 


' 

PHOSPHORUS  PENTASULPHIDE.  85 

escape.  Allow  the  apparatus  to  cool  completely,  and  remove  the 
preparation  by  the  aid  of  a  glass  rod  or  a  wire,  after  cutting  the 
tube,  if  necessary,  into  several  sections.  The  preparation  should 
be  protected  from  moisture,  best  by  keeping  it  in  a  sealed  tube. 
Yield,  slight. 

(c)  Sulphides. 

54.   Phosphorus  Pentasulphide  j1    Thermoelectric  Determination 
of  the  Boiling-Point. 

Mix  176  g.  of  flowers  of  sulphur  with  62  g.  of  dry  red  phosphorus; 
if  the  phosphorus  is  moist,  it  should  be  washed  with  hot  water, 
rinsed  on  a  filter  several  times  with  water,  then  with  alcohol,  and 
dried  in  the  hot  closet. 

Clamp  a  750  c.c.  round-bottomed  flask  to  a  ring-stand,  under  the 
hood,  at  a  suitable  height  above  a  Bunsen  burner.  Introduce  a 
spoonful  of  the  above  mixture  into  the  flask,  and  after  replacing 
approximately  all  of  the  air  with  carbon  dioxide,  heat  until  a 
reaction  begins.  Then  remove  the  burner  at  once,  and  add 
spoonful  after  spoonful  of  the  powder  so  that  each  portion  imme- 
diately enters  into  reaction;  toward  the  end  heat  the  flask  from 
time  to  time  with  the  burner.  Take  care  that  neither  the  powder 
in  the  spoon  nor  the  supply  of  unused  material  catches  fire;  keep 
a  supply  of  sand  on  hand  to  throw  on  the  blaze,  and  a  pan  to  catch 
the  liquid,  in  case  the  flask  should  break  and  its  burning  contents 
run  out.  Usually  the  preparation  can  be  carried  out  without 
accident.  After  cooling,  break  the  flask,  and  collect  the  gray  and 
somewhat  hygroscopic  crude  product. 

Purify  the  substance  by  distilling  it  from  a  small,  wide-necked, 
plain  retort,  rejecting  the  first  few  drops  of  distillate  and  collecting 
the  main  portion  in  a  dry  flask.  To  retain  the  heat,  fold  a  piece 
of  asbestos  paper  into  a  cone  and'place  it  over  the  neck  and  bulb 
of  the  retort.  The  distillate  solidifies  into  a  light-yellow,  amor- 
phous mass.  Break  the  flask  and  transfer  the  product,  after 
removing  any  fragments  of  glass,  to  a  bottle  which  can  be  tightly 
stoppered.  Yield,  about  180  g. 

A  part  of  the  preparation  can  be  further  purified  by  crystalliza- 
tion from  carbon  bisulphide.  Place  the  material  in  a  Soxhlet 
thimble,  extract  it  for  several  days  with  carbon  bisulphide  in  an 

1  Cf.  A.  Stock,  Ber.  41,  563  (1908). 


SULPHIDES. 


automatic  apparatus,  and  collect  the  crystals  which  have  separated 
in  the  boiling-flask.  The  purified  product  melts  at  275°  to  276°, 
and  0.508  part  dissolves  in  100  parts  of  boiling  carbon  bisulphide. 
Thermoelectric  Determination  of  the  Boiling-point.  Le  Chate- 
lier's  pyrometer  is  used  for  the  measurement  of  temperatures  up 
to  1600°.  It  consists  of  two  wires,  one  of  platinum  and  the  other 

of  10%  rhodium  -  platinum 
alloy,  which,  at  one  end,  are 
fused  together  in  the  oxyhy- 
drogen  flame  and  at  the  other 
end  are  connected  through  a 
suitable  measuring  instrument. 
A  difference  between  the 
temperature  at  the  junction 
of  the  wires  and  that  at  the 
free  ends  gives  rise  to  an  elec- 
tric current,  the  feeble  poten- 
tial of  which  may  be  measured 
by  means  of  a  sensitive  volt- 
meter. In  order  to  economize 
in  the  amount  of  platinum,  the 
free  ends  of  the  pyrometer  are 
connected  to  insulated  copper 
wires  leading  to  the  measuring 
instrument.  The  two  parts  of 
the  circuit  where  the  connec- 
tion with  the  copper  wires  is 
made  are  inserted  in  glass  test- 
tubes,  and  the  latter  are  closed 
by  cork  stoppers.  The  test- 
tubes  are  placed  in  a  beaker 
FIG.  18.  filled  with  water  to  serve  as  a 

thermostat,  and  the  beaker  is 

enveloped  in  asbestos  paper.  In  the  voltmeter  which  is  to  be  used 
for  pyrometric  work,  a  second  scale  is  graduated  to  give  directly  the 
temperature  readings,  each  division  of  the  scale  representing  10 
degrees.  In  order  to  find  the  true  temperature  of  the  junction,  the 
reading  must  be  corrected  by  adding  half  the  temperature  of  the 
thermostat.  If  the  latter  is  maintained  at  0°,  by  filling  the  beaker 


PHOSPHORUS  PENTASULPHIDE.  87 

with  ice,  no  correction  is  necessary.     On  account  of  the  sensitive- 
ness of  the  thermo-element  toward  chemical  influences,-  the  wires 
are  protected  by  a  long,  narrow  porcelain  tube,  to  the  closed  end  of 
which  the  thermoelectric  junction  is  inserted.     Inside 
this  tube,  one  wire  is  isolated  from  the  other  by  being 
placed  inside  a  porcelain  capillary  (a,  Fig.  19).     The 
protective  tube  should,  to  save  space,  be  made  as  nar- 
row as  possible,  and  it  may  be  of  glazed  or  unglazed 
porcelain:  the  latter  will  stand  a  higher  heat,  but  is  not 
impervious  to  gases ;  at  temperatures  up  to  red  heat,  a 
protective  tube  of  Jena  glass  may  be  employed,  which, 
particularly  if  the  wires  are  insulated  with  mica,  can 
be  made  of  smaller  bore.     In  setting  up  an  apparatus 
for   the   pyrometric    measurements    of    temperatures, 
special  care  should  be  taken  to  properly  connect  the 
thermo-element  with  the  terminals  of  the  voltmeter,      „ 
to   carefully   isolate   the  pyrometer  wires,  and  to  see 
that  the  voltmeter  is  properly  adjusted  and  set  to  the  correct 
zero  point. 

The  pyrometer  is  best  standardized  by  carrying  out  measure- 
ments with  a  few  pure  metals,  the  melting-points  of  which  are 
accurately  known.1  Refer  to  the  discussion,  under  Tin  (No.  6, 
p.  15),  of  melting-point  determinations. 

To  determine  the  boiling-point  of  the  phosphorus  penta- 
sulphide,  fasten  a  Jena-glass  test-tube,  about  40  cm.  long  and 
4  cm.  in  diameter,  in  a  perpendicular  position,  and  place  in  it 
about  40  g.  of  the  material.  Around  the  middle  part  of  the 
tube  above  the  substance,  wrap  several  layers  of  asbestos  paper 
and  secure  the  pyrometer  so  that  the  thermoelectric  junction 
hangs  two  or  three  centimeters  above  the  surface  of  the  melted 
material.  Heat  the  phosphorus  pentasulphide  to  boiling  with  a 
large  flame,  meanwhile  passing  a  slow  stream  of  carbon  dioxide 
into  the  upper  part  of  the  test-tube  in  order  to  prevent  the 
ignition  of  the  hot  vapors  when  they  come  in  contact  with  the 


1  Cf.  L.  Holborn  and  A.  Day,  Ann.  d.  Physik.  (4)  2,  545  (1900);  4,  99 
(1901). 

Cd  melts  at  321°.  Cu  melts  at  1065°  (in  air). 

Pb      "      "  327°.  Cu      "      "  1084°  (in  nitrogen). 

Zn      "      "  419°.  Au     "      "  1064°. 


88  SULPHIDES. 

air.  The  voltmeter  rises  slowly  until  it  registers  the  temperature 
of  the  hot  vapor,  after  which  it  remains  constant.  It  is  neces- 
sary, however,  that  the  liquid  should  boil  vigorously.  The 
boiling-point  determined  in  this  way  in  the  author's  laboratory 
with  the  substance  prepared  as  above,  was  507°  to  508°  at  710  mm., 
which  value  agrees  with  that  obtained  by  Bodenstein  (508°  at 
atmospheric  pressure). 

55.   Black  Mercuric  Sulphide ;  Transformation  into  Cinnabar. 

The  black  sulphide  of  mercury  is  always  formed  by  the  direct  combination 
of  the  elements,  as  well  as  by  the  precipitation  of  a  mercuric  salt  with  hydrogen 
sulphide,  but  this  modification  changes  slowly,  of  itself,  into  the  more  stable, 
red  form  (cinnabar).  The  transformation  may  be  accelerated  greatly  by 
allowing  the  black  sulphide  to  stand  in  a  warm  place  in  contact  with  a  solution 
containing  caustic  alkali  and  alkaline  sulphide;  little  by  little  the  relatively 
more  soluble  black  sulphide  dissolves,  and  the  more  insoluble  red  form 
separates  in  a  crystalline  condition. 

Triturate  50  g.  mercury,  20  g.  flowers  of  sulphur,  and  a  little 
ammonium  sulphide  solution,  in  a  large  porcelain  mortar.  Mix 
the  resulting  black  paste  of  mercuric  sulphide,  sulphur,  and 
globules  of  mercury  with  60  g.  of  a  20%  caustic  potash  solution, 
and  allow  the  mixture  to  stand  at  a  temperature  of  about  50° 
(e.g.,  on  top  of  the  hot  closet).  Replace  the  evaporated  water 
daily,  mixing  up  the  mass  each  time  with  the  pestle.  When,  at 
the  end  of  a  week  at  most,  the  mass  has  become  of  a  pure  red 
color,  wash  it  by  decantation  with  water,  whereby  the  greater 
part  of  the  excess  of  sulphur  is  removed,  and  then  decant  the 
cinnabar  itself  with  water  into  an  evaporating  dish,  allowing  any 
lumps  of  the  black  sulphide  to  remain  behind.  Boil  the  material 
with  a  sodium  sulphite  solution  to  remove  the  remainder  of  the 
free  sulphur,  and  wash  the  final  product  by  decantation  with 
boiling  water.  Drain  the  product  on  the  suction-filter  and  dry  it 
in  the  hot  closet.  Yield,  80  to  90%  of  the  theoretical. 

56.    Sulphides  of  Tin. 

Stannous  Sulphide.  Dissolve  12  g.  of  tin  in  50  g.  of  concen- 
trated hydrochloric  acid,  warming  gently.1  Dilute  the  filtered 

1  The  addition  of  2.5  c.c.  of  nitric  acid,  of  sp.  gr.  1.2,  together  with  the 
hydrochloric  acid  will  cause  the  metal  to  dissolve  very  much  more  rapidly. 
The  nitric  acid  is  reduced  to  ammonium  salt.  (Translators.) 


GREEN   MANGANESE   SULPHIDE.  89 

solution  to  2000  c.c.,  and  treat  it,  at  the  temperature  of  the  water 
bath,  with  hydrogen  sulphide  until  the  precipitate  settles  clear. 
Collect  the  stannous  sulphide  upon  a  large  plaited  filter,  wash  it 
thoroughly  with  hot  water,  and  allow  it  to  dry  over  night  on  a 
porous  plate.  Break  up  the  caked  product  which  results  and 
leave  it  in  the  drying  closet,  occasionally  breaking  up  the  lumps 
and  grinding  them  to  a  powder,  until  the  mass  is  entirely  free 
from  moisture. 

Stannic  Sulphide.  Triturate  the  stannous  sulphide,  prepared 
by  the  above  directions,  with  half  its  weight  of  sulphur  and  6  g. 
of  ammonium  chloride.  Place  the  mixture  in  an  Erlenmeyer 
flask,  which  should  be  about  three-quarters  filled,  and  bury  the 
flask  to  its  neck  on  a  sand  bath  or  place  it  on  a  Babo  funnel, 
covering  it  with  a  cone  of  asbestos  paper.  Heat  for  two  hours, 
not  too  strongly,  but  say  sufficiently  to  keep  the  iron  pan  of  the 
sand  bath  at  a  fairly  bright  red  heat.  Then,  when  the  ammo- 
nium chloride  and  the  excess  of  sulphur  should  be  completely 
volatilized,  allow  the  contents  of  the  flask  to  cool.  On  breaking 
the  flask,  the  stannic  sulphide  is  obtained  in  the  form  of  soft, 
glistening,  yellow  crystals.  It  is  known  in  this  modification 
by  the  name  of  Mosaic  Gold.  Impurities  may  be  removed 
by  decantation  with  water;  they  either  float  off  or  remain 
entirely  behind. 

Another  method  for  preparing  this  compound  is  to  start  with 
stannic  chloride  and  carry  out  the  process  as  described  for  tita- 
nium disulphide,  No.  58. 

57.   Green  Manganese  Sulphide. 

Heat  a  solution  of  40  g.  crystallized  manganous  sulphate  in 
1200  c.c.  of  water  to  boiling  in  a  2-liter  beaker.  Blow  in  a  vigor- 
ous current  of  steam,  and  add  300  c.c.  of  yellow  ammonium  sul- 
phide, all  at  one  time,  while  the  solution  is  being  agitated  by  the 
steam;  the  manganous  sulphide  formed  appears  reddish  for  the 
first  moment,  immediately  turns  yellowish,  and  soon  becomes  dark 
olive-green.  Wash  the  precipitate,  which  settles  well,  by  decan- 
tation with  boiling  water  containing  a  little  hydrogen  sulphide; 
each  volume  of  water  added  should  be  thoroughly  mixed  with  the 
precipitate  by  blowing  in  steam.  All  of  the  ammonium  salt  is 
removed  in  the  course  of  a  few  hours.  A  small  amount  of  brown 


90  SULPHIDES. 

manganese  oxide  is  formed  during  the  process,  by  oxidation  of 
the  sulphide,  but  it  floats  on  top  of  the  liquid  and  is  removed  by 
the  decantation. 

In  order  to  dry  the  preparation,  place  the  slime  of  manganous 
sulphide,  after  draining  it  as  completely  as  possible,  in  a  300  c.c. 
flask,  and  heat  it  at  100°  to  150°  in  a  small  oven  while  conducting 
hydrogen  sulphide  gas  into  the  flask.  Allow  the  dried  sul- 
phide to  cool  in  the  atmosphere  of  hydrogen  sulphide,  for  if  the 
hot  manganese  sulphide  comes  in  contact  with  air,  it  takes  fire 
and  burns.  This  spontaneous  oxidation  is  sometimes  observed 
with  the  green  sulphide,  even  when  cold.  Instead  of  treating 
the  precipitate  of  manganous  sulphide  as  just  directed,  it  may 
be  drained  rapidly  on  a  large  suction-filter,  then  washed  with 
alcohol,  and  dried  in  a  vacuum  desiccator  over  sulphuric  acid. 
Yield,  nearly  quantitative. 

58.   Titanium  Bisulphide. 

1.   From  Titanium  Tetrachloride. 

This  preparation  furnishes  an  example  of  the  decomposition  of  a  chloride 
by  hydrogen  sulphide  in  the  absence  of  water  (cf.  p.  61). 

The  apparatus  described  in  No.  53  is  used  (Fig.  17,  p.  84), 
but  instead  of  hydrogen,  hydrogen  sulphide  is  employed.  It  is 
first  passed  through  a  small  wash-bottle  containing  glycerol,  in 
order  that  the  bubbles  may  be  counted,  and  through  two  U-tubea 
containing  calcium  chloride  to  dry  the  gas.  In  the  retort  50  g- 
of  titanium  chloride  are  placed. 

First  of  all  pass  a  rapid  stream  of  hydrogen  sulphide  through 
the  apparatus  and  heat  the  combustion  tube  to  dark  redness;  then 
heat  the  titanium  chloride  in  the  retort  nearly  to  boiling,  and  keep 
it  at  that  point  by  means  of  a  small  flame.  The  vapors  of  titanium 
tetrachloride  charged  with  hydrogen  sulphide  are  decomposed  in 
the  hot  tube  to  TiS2  and  HC1:  any  unchanged  titanium  chloride 
that  collects  in  the  receiver  is  returned  to  the  retort,  again  distilled, 
and  the  process  repeated  even  a  third  time.  The  inlet  tube  for 
the  hydrogen  sulphide  must  not  be  too  narrow,  or  it  may  become 
stopped  with  titanium  disulphide  formed  within  the  retort.  When 
all  of  the  titanium  tetrachloride  has  reacted,  withdraw  the  retort, 
and,  in  order  to  remove  all  traces  of  tetrachloride  from  the  product, 
pass  a  current  of  hydrogen  or  of  carbon  dioxide  through  the  tube. 


. 
TITANIUM   BISULPHIDE.  91 

After  allowing  the  tube  to  cool,  remove  its  contents  in  the  manner 
described  in  No.  53.  A  product  is  obtained  which  consists  of 
about  10  g.  of  dark  brass-yellow  leaflets  similar  to  those  of  Mosaic 
gold. 

Analyze  one  sample  for  titanium  by  roasting  it  in  a  porcelain 
crucible  and  weighing  it  as  TiO2;  a  second  portion  for  sulphur  by 
taking  it  up  in  aqua  regia  (or  in  the  manner  described  below)  and 
precipitating  the  highly  diluted  solution  with  barium  chloride. 

2.    From  Titanium  Dioxide. 

Titanium  disulphide  was  first  obtained  by  H.  Rose  (1823)  on  conducting 
vapors  of  carbon  disulphide  over  glowing  titanium  dioxide.  The  theory  «r* 
this  method,  which  subsequently  proved  to  be  of  importance  for  the  prepa- 
ration of  sulphides,  is  the  same  as  that  of  the  preparation  of  chlorides 
from  oxides  and  carbon  tetrachloride  (cf.  Introduction  to  No.  52,  p.  81). 

Lead  a  fairly  strong  current  of  hydrogen  sulphide  through  a 
wash-bottle  containing  glycerol,  a  U-tube  filled  with  calcium 
chloride,  a  wash-bottle  containing  carbon  disulphide,  and  into  a 
piece  of  combustion  tubing  in  which  has  been  placed  4  to  5  grams 
of  finely-powdered  titanium  dioxide  (No.  52).  Heat  the  contents 
of  the  tube  to  a  bright  red  heat  by  means  of  a  row  burner  with  an 
asbestos  hood  (Fig.  3).  The  vaporization  of  the  carbon  disulphide 
may  be  hastened  by  placing  the  wash-bottle  in  a  beaker  filled  with 
lukewarm  water.  Do  not  heat  the  tube  until  the  air  has  been 
completely  expelled;  during  the  heating  revolve  it  a  little  on  its 
axis  from  time  to  time  in  order  that  new  portions  of  the  dioxide 
may  come  to  the  surface.  After  about  three  hours  the  reaction 
is  complete;  allow  the  preparation  to  cool  in  the  current  of  hydrogen 
sulphide. 

The  titanium  sulphide  has  the  appearance  of  a  brown  powder, 
but  on  being  pressed  it  assumes  a  dark-yellow,  metallic  luster. 
In  the  analysis,  titanium  may  be  determined  by  roasting,  and 
sulphur  either  by  the  method  given  above,  or  by  heating  the 
sample  in  a  current  of  oxygen  and  collecting  the  gases  evolved  in 
ammoniacal  hydrogen  peroxide  or  in  bromine  water,  eventually 
making  this  solution  acid  and  precipitating  it  with  barium  chloride. 
The  results  show  a  somewhat  too  high  titanium,  value,  and  too  low 
sulphur,  because,  when  prepared  in  this  way,  the  sulphide  invari- 
ably contains  some  titanium  dioxide. 


92  NITRIDES. 

(d)    Nitrides. 

59.   Hydrogen  Cyanide,  Mercuric  Cyanide,  Cyanogen,  and 
Dithio-oxamide. 

Carbon  combines  with  nitrogen,  absorbing  heat,  and  forms  the  nitride, 
cyanogen : 

2  C  +  2  N  +  7100  cal.  =  C2N2. 

The  endothermic  nature  of  this  substance  (cf.  p.  47)  accounts  for  the 
formation  of  cyanogen  compounds  in  the  blast  furnace,  and  for  its  presence  in 
the  sun,  as  shown  by  the  spectroscope.  When  cyanogen  is  once  formed,  its 
rate  of  decomposition  at  a  lower  temperature  is  so  small  that  it  may  be 
prepared  pure  in  an  indirect  manner  from  its  compounds  and  preserved 
undecomposed;  cf.  No.  64,  Acetylene,  and  No.  67,  Hydrogen  Peroxide. 

Free  cyanogen  gas  which  has  the  formula  C2N2  should  not  be  confused  with 
the  monovalent  radical  CN.  The  radical  CN  forms  one  of  the  most  stable 
atomic  groupings;  it  behaves  in  the  same  manner  as  the  halogen  atoms  in 
halides,  and  it  bears  the  same  relation  to  uncombined  cyanogen  as  the  chlorine 
ion,  for  example,  bears  to  free  chlorine. 

For  the  preparation  of  cyanogen  gas,  hydrogen  cyanide  (hydrocyanic  acid) 
is  first  prepared  by  Wohler's  method  of  treating  either  sodium  or  potassium 
ferrocyanide  with  sulphuric  acid.  The  hydrogen  cyanide,  together  with  the 
residues,  is  worked  over  into  mercuric  cyanide  and  the  latter  decomposed  by 
heat. 

Hydrogen    Cyanide  from    Potassium    Ferrocyanide.     Mercuric 
Cyanide. 
2  K4[Fe(CN)6]  +  3  H2SO4=  3  K2SO4+  K2Fe[Fe(CN)6]  +  6  HCN 

6  HCN  +  3  HgO  =  3  Hg(CN)2+  3  H2O. 

Place  a  400  to  500  c.c.  flask  on  a  Babo  boiling  funnel,  and 
through  the  cork,  which  fits  its  neck  tightly,  pass  a  delivery  tube 
leading  to  a  condenser;  fit  the  lower  end  of  the  condenser  by  means 
of  another  cork  stopper  into  an  adapter  that  just  dips  into 
100  c.c.  of  water  in  the  receiving  flask.  Introduce  into  the  distill- 
ing flask  100  g.  of  coarsely  broken  potassium  ferrocyanide  and  a 
cooled  mixture  of  70  g.  concentrated  sulphuric  acid  and  130  c.c. 
water.  Distil  until  the  residue  in  the  flask  consists  of  a  thin 
white  slime  ["  prussic  acid  residue/'  K2Fe[Fe(CN)6]  (?)].  Set 
aside  about  15  c.c.  of  the  distillate.  Dilute  the  remainder  with 
water  to  a  volume  of  300  to  400  c.c.  and  add  mercuric  oxide  (about 
75  grams),  shaking  well,  until  the  solution  no  longer  smells  of 
hydrocyanic  acid  (Caution)  and  shows  a  neutral  or  barely  alkaline 
reaction.  Then  pour  in  the  reserved  15  c.c.  of  hydrocyanic  acid 
and  evaporate  the  solution  under  a  good  hood  till  it  crystallizes. 
If  necessary,  purify  the  preparation  by  recrystallization.  Almost 
the  theoretical  yield  of  85  to  90  grams  is  obtained. 


CYANOGEN.  93 

In  this  and  the  following  work  regard  should  be  paid  to  the  extremely 
poisonous  nature  of  hydrocyanic  (prussic)  acid,  and  the  operations 
should  be  carried  out  under  a  well-ventilated  hood  in  the  special  room 
for  noxious  materials,  or  in  the  open  air. 

Working  up  of  the  Prussic  Acid  Residues.  Wash  the  residue 
remaining  in  the  distilling  flask  by  decantation  with  water,  and 
then  heat  it  upon  the  water  bath  with  about  100  grams  of  2-normal 
nitric  acid,  which  oxidizes  it  to  a  mass  resembling  Prussian  blue; 
the  substance  so  formed  contains  potassium  and  is  known  as 
Williamson's  Violet.  Wash  this  pigment  several  times  by  decan- 
tation  with  water,  suck  it  free  from  liquid  on  a  hardened  filter 
paper,  and  dry  it  in  the  hot  closet.  Yield,  40  grams. 

Mercuric  cyanide  can  be  prepared  from  Prussian  blue  or  William- 
son's violet  by  boiling  the  pigment  with  2.25  times  its  own  weight 
of  mercuric  oxide  and  forty  times  its  weight  of  water.  After  the 
blue  color  has  disappeared,  filter  the  solution  and  boil  the  turbid 
filtrate  with  more  of  the  Prussian  blue  until  no  more  of  it  is  de- 
colorized and  the  solution  reacts  neutral.  Add  some  animal 
charcoal,  boil  the  solution  again,  filter,  and  finally  evaporate  to 
dryness.  Recrystallize  the  residue  from  water. 

Mercuric  cyanide  in  aqueous  solution  is  but  slightly  dissociated 
electrolytically,  and  on  this  account  it  fails  to  give  a  precipitate 
when  treated  with  either  silver  nitrate  or  with  sodium  hydroxide. 
For  the  same  reason  mercuric  oxide  will  dissolve  in  solutions  con- 
taining potassium  cyanide,  forming  mercuric  cyanide  and  potas- 
sium hydroxide;  the  experiment  should  be  tried. 

Cyanogen  from  Mercuric  Cyanide.  Heat  about  one  gram  of  the 
mercuric  cyanide  strongly  in  a  dry  test-tube;  cyanogen  gas  is 
evolved  and  burns  with  a  yellowish-red  flame,  which  is  purple 
at  the  edges. 

Dithio-oxamide. 

Cyanogen  unites  with  one  or  with  two  molecules  of  hydrogen  sulphide: 
C  =  N  CSNH2 

+  H2S  =  |  (cyanthiofonnamide). 

CEEN  C  =  N 

C  =  N  CSNH, 

|  +  2  H2S  =  |  (dithio-oxamide). 

C=N  CSNH, 

Dithio-oxamide  is  sometimes  formed  in  qualitative  analysis  when  hydrogen 
sulphide  is  passed  into  the  solution  containing  potassium  cyanide  in  the  test 
for  cadmium. 


94  NITRIDES. 

1.  Saturate   100   g.   of   alcohol   with   hydrogen  sulphide   in   a 
closed  flask.     Then  conduct  into  the  liquid  some  cyanogen  gas, 
which  is  prepared  by  heating  successively  two  portions  of  10  g. 
each  of  mercuric  cyanide  in  a  test-tube  fitted  with  a  delivery 
tube.     Finally,  saturate  the  solution  once  more  with  hydrogen 
sulphide.     After  standing  for  some  time  a  mass  of  small  crystals 
is  deposited;  boil  the  mother-liquor  with  bone-black  and  evapo- 
rate  the   filtrate   in   order  to  obtain  another  crop  of   crystals. 
Recrystallize  the  product  from  alcohol,  or  from  a   mixture   of 
acetone  and  chloroform.     Yield,  3.5  to  4  g. 

2.  Treat  an  aqueous  solution  of  25  g.  copper  vitriol  with  con- 
centrated ammonia  until  the  precipitate,  which  first  forms,  just 
disappears;   decolorize  the  liquid  by  the   addition  of  a  barely 
sufficient   amount   of  potassium   cyanide   solution   (about   26   g. 
KCN)  and  then  saturate  with  hydrogen  sulphide.     The  solution 
first  becomes   yellow,  then   it   appears   red,  and   finally  a  red, 
crystalline  powder  separates.     After  standing  for  several  hours  in 
the  ice-chest,  separate  the  crystals  from  the  liquor  and  recrystal- 
lize  them  from  a  little  alcohol.     The  yield  is  small  (about  0.5  g.) 
because  in  aqueous  solution-  the  greater  part  of  the  cyanogen  is 
reduced  to  hydrocyanic  acid.     This  is  the  reason  for  the  use  of 
alcohol  in  the  first  method. 

60.   Boron  Nitride,  BN.1 

Heat  an  intimate  mixture  of  5  g.  finely  powdered,  anhydrous 
borax  and  10  g.  ammonium  chloride,  as  hot  as  possible  with  the 
blast  lamp  in  a  covered  platinum  crucible  surrounded  by  a  clay 
mantle.  After  cooling,  pulverize  the  porous  contents  of  the 
crucible,  and  extract  it  several  times  with  water  containing  a 
little  hydrochloric  acid.  Boil  the  residue  with  pure  water,  collect 
it  on  a  filter  and  dry  it.  Yield,  about  0.3  g.  Double  this  yield 
may  be  obtained  if  a  long  platinum  crucible  ("  finger  crucible  ") 
is  available,  which  is  about  10.5  cm.  high,  3  cm.  wide  at  the  top,, 
and  1  cm.  at  the  bottom.  Place  a  layer  of  1  g.  ammonium  chloride 
in  the  bottom,  then  add  the  same  mixture  as  before.  Suspend 
the  crucible  in  an  almost  horizontal  position  by  means  of  two 
rings  made  of  iron  wire.  Slip  a  piece  of  asbestos  board,  with  a 


1  Cf.  Stock  and  Holle,  Ber.  41,  2095,  regarding  the  preparation  of  pure  boron, 
nitride. 


MAGNESIUM  NITRIDE.  95 

hole  of  the  proper  size,  a  little  way  over  the  lower  end  of  the 
crucible  in  order  that  this  end  may  be  kept  cooler.  Heat  the 
middle  part  of  the  crucible  to  bright  redness  by  means  of  a 
powerful  burner,  and  finally,  at  the  end  of  the  process,  heat  the 
bottom  with  a  Bunsen  flame  so  that  ammonium  chloride  vapor 
is  driven  through  the  glowing  mass.  The  experiment  is  finished 
when  all  of  the  ammonium  chloride  has  been  volatilized. 

Boron  nitride  is  insoluble  in  water  and  in  acids;  on  being 
boiled  with  caustic  soda  solution,  it  decomposes  slowly,  evolving 
ammonia;  it  is  attacked  more  rapidly  by  fusion  with  sodium 
carbonate  on  platinum  foil. 

61.   Magnesium  Nitride,  Mg3N2;    Ammonia  from  the  Atmosphere. 

Fill  a  small  iron  crucible  (3  cm.  high  and  5  cm.  wide  at  the 
top)  to  two-thirds  with  8  to  9  g.  of  magnesium  powder,  and  make 
a  tight  joint  with  the  cover  by  means  of  wet  asbestos  pulp. 
Likewise  close  with  asbestos  pulp  a  small  hole  which  has  been 
made  in  the  cover.  After  drying,  make  a  perforation  in  this  last 
mass  of  asbestos  by  means  of  a  fine  needle.  Dry  the  crucible, 
with  its  contents,  in  the  hot  closet. 

Place  the  crucible  in  a  hole  made  in  a  piece  of  heavy  asbestos 
board,  so  that  its  greater  part  hangs  below  the  asbestos,  and 
direct  the  flame  of  a  blast  lamp  side  wise  against  it.  By  this 
arrangement  the  flame  gases  are  kept  well  away  from  the  cover 
of  the  crucible.  Turn  the  asbestos  board  with  the  crucible 
from  time  to  time  so  that  all  sides  are  heated  equally.  Con- 
tinue the  heating  for  30  minutes. 

After  cooling,  remove  the  cover  of  the  crucible  and  the  upper 
layer  of  white  magnesium  oxide.  Beneath  the  latter  lies  a  light 
yellowish-green  mass  of  nearly  pure  magnesium  nitride,  which  is 
obtained  to  about  80%  of  the  theoretical  yield.  If  the  heating  is 
not  continued  long  enough,  dark  places  are  found  in  the  mass, 
caused  by  the  presence  of  unchanged  magnesium. 

Magnesium  nitride  reacts  readily  with  water  with  the  forma- 
tion of  ammonia.  To  obtain  the  ammonia,  place  the  whole  con- 
tents of  the  crucible  (without  separating  the  oxide  from  the 
nitride)  in  a  round-bottomed  flask  which  is  provided  with  a 
separatory  funnel  and  with  a  right-angled  tube  leading  to  a  con- 
denser. Prolong  the  condenser  with  an  adapter  that  just  dips 


96  NITRIDES. 

into  100  c.c.  of  water  in  a  small  flask.  Cool  the  evolution  flask 
by  placing  it  in  a  dish  of  cold  water,  and  then,  at  first  slowly  and 
later  more  rapidly,  allow  75  c.c.  of  water  to  drop  upon  the 
nitride.  Finally,  heat  the  contents  of  the  flask  with  a  small 
flame  as  long  as  ammonia  continues  to  be  evolved.  Yield,  3  to 
4  g.  of  ammonia,  as  determined  by  titrating  an  aliquot  part  of 
the  distillate  with  0.5  normal  acid. 

Calcium  Nitride.  Place  a  small  shaving  of  metallic  calcium 
on  an  inverted  porcelain  cover  and  set  fire  to  it  with  a  match. 
It  burns  with  a  light,  yellowish-red  flame,  and  both  oxide  and 
nitride  are  formed.  Moisten  the  white  combustion  product 
with  a  few  drops  of  water,  and  test  with  Nessler's  reagent  for 
ammonia. 

The  chemical  reaction  is  the  same  for  both  the  magnesium  and 
calcium  nitrides.  A  part  of  the  metal  serves  to  free  the  air  of 
oxygen,  while  the  remainder  combines  with  the  nitrogen  that  is 
left.  By  the  hydrolysis  of  the  nitride,  ammonia  and  the  hydrox- 
ide of  the  metal  are  formed.  The  possibility  of  obtaining 
ammonia  from  the  nitrogen  of  the  air,  with  the  intermediate  for- 
mation of  a  nitride,  was  first  pointed  out  by  Wohler  in  1850 
when  his  method  of  preparing  boron  nitride  was  published. 

63.    Chromium  Nitride,  CrN. 

Heat  5  to  10  g.  of  violet,  anhydrous  chromium  chloride  (No.  44) 
in  a  25  to  50  cm.  tube  of  difficultly  fusible  glass,  at  first  gently 
and  then  strongly  with  a  row  burner  (Fig.  3).  Meanwhile  pass 
through  the  tube  a  current  of  ammonia  gas,  which  is  obtained  by 
heating  a  concentrated  solution  of  ammonia  and  drying  the  gas 
successively  in  a  lime-tower  and  a  U-tube  containing  lime.  Leave 
the  reaction  tube  entirely  open  at  one  end,  since  if  an  ordinary 
delivery  tube  is  used  to  carry  away  the  waste  gases,  it  soon  becomes 
stopped  with  sublimed  ammonium  chloride.  Carry  out  the 
experiment  under  the  hood  and  continue  the  heating  until  no  more 
vapors  of  ammonium  chloride  escape.  After  cooling,  pulverize 
the  reaction  product  and  once  more  ignite  it  in  an  atmosphere  of 
ammonia.  Yield,  nearly  theoretical. 

Chromium  nitride  is  very  stable  towards  warm  caustic  soda 
solution,  and  hot  concentrated  sulphuric  acid  reacts  with  it  but 
slowly.  To  test  its  purity,  boil  a  small  sample  with  chlorine-free 


CALCIUM  CARBIDE.  97 

caustic  soda  and  test  the  solution  obtained  for  chloride.  If  it  is 
desired  to  remove  traces  of  chromium  chloride  from  the  prepara- 
tion, it  may  be  treated  in  the  cold  with  dilute  hydrochloric  acid  and 
a  little  tinfoil,  then  washed  with  water,  drained,  and  dried  at  110° 
to  120°. 

Chromium  nitride,  on  being  heated  upon  a  porcelain  crucible 
cover  over  the  blast  lamp,  changes  to  dull,  grayish-green  chromic 
oxide. 

(e)   Phosphides. 
63.   Magnesium  Phosphide,  Mg3P2. 

Draw  out  one  end  of  a  35  cm.  long  combustion  tube  so  that  its 
diameter  is  0.6  to  0.9  cm.,  and  connect  this  end  by  a  piece  of 
rubber  tubing  with  two  drying  bottles  and  a  Kipp  hydrogen 
generator.  Insert  two  porcelain  boats  into  the  tube,  the  one 
nearest  to  the  source  of  hydrogen  containing  about  five  grams  of 
red  phosphorus  and  the  other  four  grams  of  magnesium  powder. 
The  phosphorus  must  be  perfectly  dry  (cf.  p.  88).  Erect  the 
apparatus  under  the  hood. 

After  filling  the  apparatus  with  hydrogen  (Test!),  heat  the  tube 
slightly  through  its  entire  length  to  expel  moisture;  then,  while 
maintaining  a  steady  current  of  hydrogen,  heat  both  boats  as 
equally  as  possible  by  means  of  two  Bunsen  burners  so  that  phos- 
phorus vapor  is  carried  over  the  magnesium  and  a  violent  reaction 
takes  place.  When  all  the  phosphorus  has  volatilized,  allow  the 
contents  of  the  tube  to  cool  in  the  atmosphere  of  hydrogen. 

Magnesium  phosphide  when  thrown  into  water  is  decomposed, 
and  phosphine  is  set  free.  This  gas  has  a  very  offensive  odor;  it 
is  combustible,  but  it  does  not,  like  the  impure  phosphine  that  is 
commonly  prepared,  take  fire  spontaneously. 

(/)   Carbides. 

64.    Calcium  Carbide;  Acetylene  from  Calcium  Carbide; 
Benzene  from  Acetylene. 

Calcium  Carbide.  Heat  0.5  g.  of  lampblack  strongly  in  a 
.porcelain  crucible  for  a  few  minutes;  after  it  has  cooled  mix  with 
it  0.5  g.  of  thin  shavings  of  metallic  calcium,  and  heat  the  mixture 
strongly  with  the  blast  lamp  for  a  few  minutes.  After  cooling 


98  CARBIDES. 

there  is  found  in  addition  to  lampblack  a  sintered  white  mass 
which  contains  some  calcium  carbide;  on  testing  it  with  water, 
acetylene  is  formed. 

Acetylene  from  Calcium  Carbide.  Allow  water  to  drop  slowly 
from  a  dropping  funnel  upon  a  few  pieces  of  commercial  calcium 
carbide  in  a  half-liter  flask.  Pass  the  acetylene  evolved  through 
a  solution  of  sodium  plumbite  in  order  to  remove  any  hydrogen 
sulphide,  and  through  an  acid  solution  of  copper  sulphate  to  take 
out  any  phosphine.  Acetylene  burns  with  a  brilliantly  luminous 
and,  unless  in  a  special  burner,  very  smoky  flame.  Before  ignit- 
ing the  gas  it  should  be  tested  to  see  whether  all  the  air  has  been 
removed  from  the  apparatus. 

At  the  temperature  of  the  electric  arc,  acetylene  can  be  formed  from  its 
elements  with  absorption  of  heat.  At  a  bright  red  heat  it  decomposes,  but  at 
lower  temperatures,  although  it  is  likewise  unstable,  its  rate  of  decomposition 
is  extremely  slow.  (Compare  with  the  similar  relations  which  exist  for 
cyanogen,  No.  59,  and  hydrogen  peroxide,  No.  67.) 

At  a  dull  red  heat  acetylene  polymerizes  to  a  considerable  extent,  forming 
benzene:  3  C2H2  =  C6H6.  The  presence  of  benzene  in  coal-tar  may  be 
attributed  to  this  reaction ;  indeed,  this  supposition  is  supported  by  the  fact 
that  the  yield  of  benzene  in  the  tar  sinks  if  the  temperature  of  the  gas  retorts 
is  raised. 

Benzene  from  Acetylene.  Prepare  pure  acetylene,  as  directed 
above,  and  pass  a  strong  current  of  it  for  two  hours  through  a 
glass  tube  that  is  heated  to  dull  redness  in  a  combustion  furnace 
(the  air  must  be  expelled  by  acetylene  before  the  tube  is  heated  !). 
Lead  the  products,  escaping  from  the  combustion  tube,  through  a 
condenser;  in  this  way  a  few  cubic  centimeters  of  an  oil  having 
the  odor  of  coal-tar  are  obtained.  By  distilling  this  product, 
a  crude,  colorless  benzene  is  obtained  which  boils  between  70° 
and  90°.  For  the  further  identification  of  this  substance  as  ben- 
zene it  may  be  converted  into  nitrobenzene,  aniline,  and  then 
mauvein.  Refer  to  a  text-book  on  organic  preparations,  for  ex- 
ample Gattermann,  "Practical  Methods  of  Organic  Chemistry" 
(translated  by  W.  S.  Shober). 


CHAPTER  IV. 


COMPOUNDS   CONTAINING   A   COMPLEX   NEGATIVE 
COMPONENT. 

THE  modern  conception  of  complex  compounds  arose  from  the  necessity  of 
classifying  the  compounds  formed  by  the  union  of  two  simple  salts.  Those 
composite  salts  which  in  aqueous  solution  are  dissociated  entirely  into  the 
simple  salts,  or  their  ions,  are  classed  as  double  salts;  those  which,  instead  of 
dissociating  into  the  simple  salts,  give  characteristic  ions  of  their  own  —  for 
example,  simple  metal  cations  and  composite  metal-containing  anions  —  are 
classed  as  complex  salts: 

KMgCl3  =  KC1   +  MgCl2  =  K+  +  Mg++  +  3  Cl~  (double  salt), 
K2[HgIJ  =  2  K++  [HglJ™  (complex  salt). 

In  such  a  complex  compound,  the  composite  part  which  remains  intact 
during  dissociation  is  known  as  a  complex  radical  or  complex  ion  l  and  may  be 
regarded  as  the  negative  component  of  a  new  acid ;  sometimes  indeed  the  com- 
plex acid  can  be  prepared  in  a  free  state,  for  example,  hydroferrocyanic  acid 
(No.  108).  The  individual  constituents  of  a  complex  ion,  since  they  are  not 
present  in  the  free  state,  do  not  show  all  the  reactions  which  are  characteristic 
of  the  simple  ions;  it  is  this  reduced  ability  of  metals  contained  in  complex 
radicals  to  react,  which  forms  the  most  important  criterion  of  this  class  of 
compounds. 

The  conception  of  complex  compounds  has,  however,  extended  considerably 
beyond  the  range  of  these  composite  salts.  In  the  first  place,  the  distinction 
between  double  and  complex  salts  is  merely  a  qualitative  one.  Transition 
forms  are  known  in  which  the  simple  salts  and  their  ions  as  well  as  the  complex 
ions  are  found  among  the  dissociation  products;  the  nature  and  extent  of  the 
dissociation  is  also  dependent  to  a  marked  degree  upon  the  nature  of  the  solvent 
and  upon  the  concentration.  Again,  no  essential  distinction  can  be  drawn 
between  the  compounds  with  complex,  metal-containing  anions  and  those  which 
are  formed  by  the  addition  of  NH3,  H2O,  NO2,  etc.,  to  the  metal  ion  of  a  simple 
salt,  and  in  which  the  metal-containing  cation  is  complex  and  the  activity  of 
its  constituents  is  restricted.  Finally,  certain  binary  substances,  such  as  for 
example  the  oxides,  which  themselves  possess  none  of  the  characteristics  of 
salts,  can  produce  salts  (or  acids  or  bases)  by  uniting  with  one  another.  Thus 


1  It  is  a  quite  common  practice  to  indicate  complex  ions  in  chemical  for- 
mulas by  enclosing  them  in  brackets. 

99 


100  COMPLEX  COMPOUNDS. 

the  ordinary  oxygen-acids  are  formed  by  the  union  of  water  and  acid  anhy- 
dride, and  the  salts  of  the  oxygen-acids  by  the  union  of  metal  oxide  and 
non-metal  oxide.  Here  also  there  are  produced,  quite  in  accord  with  the 
interpretation  of  complex  salts,  new  ions,  the  separate  constituents  of  wrhich 
are  incapable  of  entering  into  independent  chemical  reactions: 

H2O  +  SO3  =  HJSOJ, 

H2[SOJ  =  2^  +  180,]—. 

Thus  the  sulphate  radical  does  not  show  the  reactions  of  sulphur  any  more  than 
the  ferrocyanide  radical  exhibits  those  of  iron. 

Since  manifestly  every  compound  containing  more  than  two  elements  can 
under  certain  conditions  behave  in  such  a  manner  as  to  indicate  the  grouping 
of  two  of  its  constituents  into  a  complex,  it  appears  more  rational  to  include 
under  complex  compounds  all  substances  which  are  produced  from  simple  com- 
pounds by  the  addition  of  one  or  several  elements;  in  most  cases  the  complex 
radical  behaves,  like  the  hydroxyl  group,  as  a  unit,  and  can  thus  be  treated  as 
a  substituent  of  an  atom. 

The  most  essential  difference  between  the  complex  and  the  simple  com- 
pounds lies  in  the  great  variety  of  ways  in  which  the  former  can  react  or  dis- 
sociate, and  this  is  more  evident  in  proportion  to  the  number  of  constituents. 
There  is  but  one  way  in  which  electrolytic  and  non-electrolytic  dissociation  can 
take  place  in  simple  compounds.  With  complex  compounds,  on  the  other 
hand,  the  point  of  division  is  usually  not  the  same  by  the  electrolytic  as  by  the 
non-electrolytic  dissociation: 

CaCO3  =  Ca        +  CO3  (electrolytic  dissociation).1 

CaCO3  =«  CaO     +  CO2  (non-electrolytic  dissociation).1 

Non-electrolytic  dissociation  may  be  of  very  different  types: 

KC1O3  =  KC1  +  3  O  (Nos.  74  and  75). 

NaNO3  =  NaNO2  +  O  .  (No.  79). 

Na2S2O3  =  Na2SO3  +  S  (No.  90). 

KPbI3  =  KI  +  PbI2  (No.  105). 

NH4C1  =  NH3  +  HC1  (No.  120). 

Ni(NH3)6Br2  =  NiBr2  +  6  NH3  (No.  128). 

CLASSIFICATION  OF  COMPLEX  COMPOUNDS.  Of  the  various  ways  in  which 
complex  compounds  can  dissociate,  the  electrolytic  dissociation  exceeds  all  others 
as  regards  frequency  of  occurrence.  We  may  classify  the  complex  compounds, 
therefore,  as:  1.  Those  with  a  complex  radical  which  yields  an  ion  of  negative 
charge  (Chapter  IV);  2.  Those  with  a  complex  radical  which  yields  an  ion  of 
positive  charge  (Chapter  V);  3.  Those  characterized  by  very  little  or  no 
capacity  for  electrolytic  dissociation  (Chapter  VI). 


1  According  to  whether  the  one  or  the  other  possibility  of  dissociation  is  to 
be  brought  especially  to  notice,  two  methods  of  writing  the  name  and  symbol 
have  been  devised:  CO2  •  CaO,  "  carbonate  of  lime  ";  CaCO3,  "  calcium  car- 
bonate ";  but  it  would  be  just  as  biased  to  defend  the  older  name  as  being  the 
only  satisfactory  one  as  it  would  be  to  reject  it  as  "  unscientific." 


SODIUM  PEROXIDE.  101 

Within  the  class  of  compounds  possessing  complex  anions  that  sub-class 
will  first  be  treated  which  comprises  the  compounds  containing  homogeneous 
complexes.  Homogeneous  complex  anions  arise,  as  the  name  implies,  when  one 
or  more  additional  atoms  of  the  same  element  are  joined  to  the  negative  con- 
stituent of  a  binary  compound, 

K2S  +  3  S  =  K2S4. 

Homogeneous  complex  compounds  can  be  distinguished  from  simple  com- 
pounds of  the  general  formula  Am  Bn  by  the  fact  that  the  greater  number  of 
negative  atoms  is  not  due  to  an  increased  valence  of  the  positive  component, 
but  rather  to  an  increased  combining  power  of  the  negative  part  originally 
present  in  the  mother-substance.  FeCl2  differs  from  FeCl3  in  the  valence  of 
the  iron.  Aqueous  solutions  of  FeCl2  and  FeCl3  are  identical  with  regard  to 
their  anions  but  different  with  regard  to  the  cations;  in  aqueous  solutions  of  KI 
and  KI3,  however,  the  exact  opposite  is  true. 

Homogeneous  complex  cations  are  not  known. 

The  part  played  by  the  valence  theory  in  the  interpretation  of  complex 
compounds  is  varied.  In  some  cases,  as  with  sulphuric  acid,  it  offers  certain 
advantages  in  its  original  form.  In  other  cases  an  enlargement  of  our  valence 
ideas  seems  necessary  (see  Complex  Halogen  Salts,  p.  139). 

COMPOUNDS    WITH    HOMOGENEOUS    COMPLEXES. 

(a)    Peroxides. 
65.    Sodium  Peroxide. 

Arrange  a  train  of  apparatus  so  that  air  may  be  drawn  by 
means  of  a  suction  pump  through  one  wash-bottle  containing 
sodium  hydroxide  and  two  more  containing  concentrated  sul- 
phuric acid,  and  then  into  a  wide,  30  cm.  long,  combustion  tube 
which  is  to  be  heated  on  a  row  burner  with  an  asbestos  cover 
(Fig.  3).  The  other  end  of  the  combustion  tube  is  made  nar- 
rower and  fitted  to  an  air  filter  which  serves  to  hold  back  sodium 
peroxide  dust.  This  filter  consists  of  a  30  cm.  long  and  3  cm. 
wide  glass  tube  loosely  filled  with  asbestos  fibers.  Between 
the  air  filter  and  the  pump,  place  an  empty  suction  flask  to  serve 
as  a  safety  bottle.  Into  the  combustion  tube  introduce  an 
aluminium  boat,  about  16  cm.  long  and  as  deep  as  possible, 
which  can  be  prepared  from  thin  aluminium  foil;  in  the  boat  place 
2  to  3  g.  of  sodium.  Since,  during  the  progress  of  the  experi- 
ment, the  sodium  will  tend  to  flow  in  a  direction  opposite  to 
that  of  the  air  current,  the  combustion  tube  together  with  the 
burner  should  be  placed  in  a  slanting  position  so  that  the  end 
at  which  the  air  enters  is  uppermost. 


102  PEROXIDES. 

Heat  the  sodium  to  above  its  melting-point  and  draw  air  over 
it;  at  about  300°,  it  takes  fire.  Then  lower  the  flame,  for 
from  that  point  on  the  combustion  proceeds  almost  without  any 
application  of  external  heat.  Continue  to  pass  a  vigorous  cur- 
rent of  air  until  the  reaction  has  ceased.  The  yield  is  4  to  5  g. 
of  bright  yellow  sodium  peroxide,  which  must  be  protected  from 
moisture. 

66.   Barium  Peroxide. 

At  a  dull  red  heat,  dry  barium  oxide  is  changed  by  the  oxygen  of  the  air  to 
barium  peroxide.  In  carrying  out  the  reaction,  the  air  must  be  free  from 
moisture  and  from  carbon  dioxide.  Since  the  barium  peroxide  gives  off  its 
oxygen  again  at  the  same  temperature  and  a  reduced  oxygen  pressure,  or  at 
the  same  pressure  and  a  higher  temperature,  a  technical  process  for  obtaining 
pure  oxygen  from  the  air  according  to  the  reversible  reaction 

BaO  +  O  «=>  BaO, 

has  been  developed.  Barium  oxide  cannot  be  prepared  conveniently  from 
barium  carbonate,  at  least  not  on  a  small  scale,  because  the  decomposition 
temperature  of  the  latter  lies  too  high. 

Crude  Barium  Peroxide.  Place  130  g.  of  barium  nitrate  in  a 
large  clay  crucible  and  heat  it  in  a  charcoal  furnace  with  slowly 
rising  temperature,  to  a  dull  red  heat.  After  cooling  break 
the  hard,  porous,  gray  contents  of  the  crucible  into  small 
lumps  and  transfer  it,  before  it  can  attract  moisture,  into  a 
weighed  combustion  tube.  Determine  the  quantity  of  barium 
oxide  by  weighing  the  tube  and  contents.  Heat  the  tube  in  a 
combustion  furnace  to  a  barely  perceptible  redness,  and  then 
draw  through  it  a  current  of  air  which  has  passed  through  caustic 
soda  solution  and  then  through  concentrated  sulphuric  acid.  After 
two  or  three  hours,  let  the  tube  cool  and  determine  the  gain  in 
weight,  which  should  be  one-tenth  the  original  weight  of  the 
barium  oxide;  if  it  is  less,  the  preparation  must  be  heated  again 
in  the  current  of  air. 

Pure  Barium  Peroxide.  Dissolve  the  barium  peroxide,  a 
little  at  a  time,  in  the  calculated  amount  of  ice-cold  1%  hydro- 
chloric acid  and  add  barium  hydroxide  solution  to  the  cloudy 
liquid  until  a  precipitate  consisting  chiefly  of  metal  hydroxides, 
which  are  present  as  impurities,  just  forms.  Filter,  and  precipi- 
tate the  filtrate  completely  with  barium  hydroxide,  whereby  a  fine, 
crystalline  powder  of  barium-peroxide-hydrate,  BaO2  •  8  H20,  is 
formed.  Drain  the  precipitate  and  dry  it  in  the  steam  closet. 


HYDROGEN   PEROXIDE.  103 

67.   Hydrogen  Peroxide. 

The  oxidation  of  water  to  hydrogen  peroxide  takes  place  with  absorption  of 
heat: 

2  H2O  (liquid)  +  O2  +  44,320  cal.  =  2  H2O2  (liquid). 

Hydrogen  peroxide  is  thus  endothermic  as  regards  its  formation  from  water 
and  oxygen.  In  accordance  with  the  relation  which  exists  between  the  equi- 
librium constants  of  the  mass-action  law  and  the  temperature  (cf.  p.  47),  it 
follows  that  the  quantity  of  hydrogen  peroxide  in  an  equilibrium  mixture  must 
increase  with  rise  of  temperature,  and  that  therefore  in  order  to  obtain  hydro- 
gen peroxide  synthetically  by  the  above  reaction  it  is  necessary  to  work  at  as 
high  a  temperature  as  possible.  Even  at  2000°  there  is  but  little  peroxide  pres- 
ent in  the  equilibrium  mixture.1  At  low  temperatures  hydrogen  peroxide  is 
not  stable  at  any  appreciable  concentration;  in  fact  a  liquid  containing  a  high 
percentage  of  it  is  explosive.  The  velocity  at  which  hydrogen  peroxide  decom- 
poses, when  in  the  region  of  instability  as  regards  temperature  or  concentration, 
increases  (as  does  the  velocity  of  all  reactions)  with  rise  of  tempera- 
ture.2 If  then  it  is  desired  to  obtain  the  hydrogen  peroxide  produced  by  a 
reaction  at  a  high  temperature,  it  is  necessary  to  cool  the  reaction  products 
very  rapidly  to  a  point  where  the  decomposition  velocity  is  inappreciable. 
When  thus  chilled  the  hydrogen  peroxide  continues  to  exist,  as  it  were,  in  a 
supercooled  condition.  This  end  is  accomplished,  for  example,  when  an  oxy- 
hydrogen  flame  comes  in  direct  contact  with  a  piece  of  ice  and  thereby  an 
extremely  sudden  drop  in  temperature  is  brought  about.  Cf.  Cyanogen, 
No.  59,  and  Acetylene,  No.  64. 

The  principle  of  preserving  the  equilibrium  concentrations,  as  they  exist  at 
high  temperatures,  by  means  of  sudden  cooling  was  first  introduced  by  Deville 
in  1863  and  applied  in  the  construction  of  his  "  hot  and  cold  tubes."  A  nar- 
row silver  tube  cooled  by  running  water  was  placed  in  the  center  of  a  white-hot 
porcelain  tube  at  the  walls  of  which  the  gas-reaction  to  be  measured  reached  its 
high-temperature  equilibrium;  the  reaction-products  on  coming  in  contact 
with  the  inner  tube  were  chilled,  and  thus  prevented,  partially  at  least,  from 
undergoing  the  reverse  reaction.  In  recent  years  many  equilibria,  in  which 
concentrations  of  measurable  magnitude  are  reached  only  at  high  tempera- 
tures, have  been  studied  in  this  manner.  The  mixtures  are  cooled,  without 
suffering  change  in  concentration,  to  temperatures  at  which  analytical  meas- 
urements are  possible.  The  quantitative  preservation  of  the  concentration 
fails,  however,  in  the  case  of  hydrogen  peroxide  on  account  of  its  great  decom- 
position-velocity. 

The  velocity  of  many  reactions  can  be  increased  by  means  of  catalyzers,  as 
well  as  by  rise  of  temperature.  The  decomposition  of  hydrogen  peroxide, 
since  it  takes  place  with  conveniently  measurable  rapidity  at  ordinary  tem- 

1  At  2000°  according  to  Nernst  (1903)  less  than  1%  H2O2  exists  in  a  mixture 
of  water  vapor  and  oxygen,  each  at  0.1  atmosphere  pressure. 

2  In  order  to  keep  the  temperature  low  during  the  distillation  of  hydrogen 
peroxide,  it  is  customary  to  work  in  a  vacuum. 


104  PEROXIDES. 

peratures,  offers  an  excellent  opportunity  for  studying  the  effect  of  various 
catalyzers.  In  this  way  a  remarkable  analogy  between  organic  ferments  and 
inorganic  catalyzers  has  been  discovered.  Cf .  No.  20.  Hot  solutions  of  hydro- 
gen peroxide  are  decomposed  rapidly  at  the  rough  places  on  porcelain  or  glass 
apparatus. 

A.  Hydrogen  Peroxide  from  the  Oxy-hydrogen   Flame.     Let  a 
hydrogen  flame  one-half  centimeter  long,  burning  from  a  glass 
tip,  play  against  a  piece  of  ice  in  a  watch  glass,  until  all  the  ice 
is  melted.     The  presence  of  hydrogen  peroxide  in  the  resulting 
liquid  may  be  easily  shown  by  means  of  a  titanium  solution 
(seep.  83). 

B.  Hydrogen  Peroxide  from  Barium  Peroxide.     Add,  little  by 
little,  some  pulverized,  crude  barium  peroxide1   (No.   66)  to  a 
mixture  of  0.6  of  its  weight  of  concentrated  sulphuric  acid  and 
300  c.c.  of  water;  the  liquid  must  be   kept   cold   by  ice  both 
inside  and  outside  the  vessel.     Neutralize  the  excess  of  acid  with 
barium  carbonate,  let  the  precipitate  settle,  filter,  and  distil  the 
solution  on  the  water  bath. 

For  the  distillation  use  a  liter  round-bottomed  flask  with  a 
stopper  through  which  pass  two  glass  tubes;  one  is  quite  narrow 
and  terminates  below  in  a  long,  fine  capillary  and  above  in  a 
rubber  tube  with  a  screw-cock.  The  other  tube  leads  from  the 
flask  to  a  Jong  condenser;  the  lower  end  of  the  latter  is  fitted  to 
a  suction  bottle  which  serves  as  a  receiver.  All  parts  of  the 
apparatus  must  be  closed  with  tightly-fitting  stoppers.  Evacuate 
the  apparatus,  but  allow  a  fine  stream  of  air  bubbles  to  flow  con- 
tinuously from  the  capillary  through  the  liquid  in  the  flask,  in 
order  to  avoid  bumping;  the  rate  of  flow  of  this  air  current  is  to 
be  regulated  by  the  screw-cock.  Distil  over  about  one-third  of 
the  contents  of  the  flask  and  test  the  distillate  for  hydrogen  per- 
oxide as  described  below;  it  should  be  practically  pure  water. 
Then  empty  the  receiving  bottle  and  distil  as  before  until  all  of 
the  remaining  liquid  has  passed  over. 

Qualitative  Tests  for  Hydrogen  Peroxide. 

1.  Titanium  Sulphate  Test  (cf.  p.  83). 

2.  Chromic  Acid  Test:  To  a  few  c.c.  of  a  chromate  solution 
weakly  acidified  with  sulphuric  acid,  add  a  few  drops  of  hydrogen 


The  use  of  barium  peroxide  hydrate  has  also  been  recommended. 


HYDROGEN  PEROXIDE.  105 

peroxide,  and  then  shake  with  about  1  c.c.  of  ether;  an  intensely 
blue  compound  of  chromium,  is  formed  which  is  more  soluble 
in  ether  than  in  water  and  passes,  therefore,  into  the  ether  layer. 

3.  A  very  dilute  solution  of  potassium  iodide  on  being  treated 
with  a  small  amount  of  hydrogen  peroxide  slowly  turns  yellow  (or 
blue  if  starch  is  also  added)  as  a  result  of  the  separation  of  iodine. 
This  reaction  is  catalytically  accelerated,  somewhat  by  acetic  acid, 
more  by  mineral  acids,  and  most  of  all  by  ferrous  sulphate. 

Quantitative  Determination  of  Hydrogen  Peroxide.  Dilute 
10  c.c.  of  the  preparation  to  about  300  c.c.,  acidify  strongly  with 
sulphuric  acid,  and  titrate  with  potassium  permanganate: 
2  KMnO4  +  3  H2SO4  +  5  H2O2  =  K2SO4  +  2  MnSO4  +  8  H2O  +  5  O2. 
1  c.c.  of  0.1-normal  KMnO4  solution  =  0.0017  g.  H2O2. 

C.  Hydrogen  Peroxide  as  a  By-product  in  Slow  Atmospheric 
Oxidations. 

Schoenbein,  in  1864,  established  the  fact  that  hydrogen  peroxide  is  produced 
when  metals,  or  their  amalgams,  are  shaken  with  water  containing  dissolved 
oxygen: 

Zn  -+  2  H2O  +  O2  =  Zn(OH)2  +  H2O2. 

It  was  shown  by  M.  Traube,1  in  1893,  that  the  formation  of  hydrogen  per- 
oxide according  to  the  equation  becomes  quantitative  if  the  reaction  is  allowed 
to  take  place  in  the  presence  of  calcium  hydroxide,  for  then  the  peroxide  is 
precipitated  as  the  difficultly  soluble  calcium  peroxide  and  is  thus  withdrawn 
from  the  sphere  of  action. 

Such  an  oxidation  in  which  twice  as  much  oxygen  is  used  as  is  necessary  for 
the  primary  process  itself,  the  other  half  of  the  oxygen  being  used  in  some 
secondary  reaction  —  mostly  in  the  formation  of  hydrogen  peroxide  —  is  of 
frequent  occurrence  both  in  organic  and  inorganic  chemistry  and  is  known  as 
'*  auto-oxidation."  Inasmuch  as  both  the  primary  and  secondary  reactions 
require  the  same  quantity  of  oxygen  and  since  they  do  not  take  place  inde- 
pendently of  one  another,  —  hydrogen  peroxide  is  not  formed  at  ordinary  tem- 
peratures from  water  and  oxygen  (see  above),  —  it  is  necessary  to  assume  the 
existence  of  some  peculiar  form  of  reaction-mechanism.  Views  regarding  this, 
however,  differ  widely ;  for  example,  the  hydrogen  peroxide  may  be  regarded 
as  a  decomposition  product  of  some  higher  oxide,  such  as  zinc  peroxide,  which 
is  formed  in  the  primary  reaction ;  or  it  may  be  assumed  with  Traube  that  a 
dissociation  of  the  water  molecules  takes  place,  the  hydroxyl  groups  combining 
with  the  metal  and  the  hydrogen  atoms  combining  with  undissociated  oxygen 
molecules. 

Place  200  c.c.  of  water,  6  g.  of  slacked  lime,  2.5  g.  of  potas- 
sium hydroxide,  and  30  g.  of  mercury  in  a  thick-walled  liter 

1  Ber.  26,  1471  (1893). 


106  POLYSULPHIDES. 

flask.  Amalgamate  the  surface  of  about  0.5  g.  of  zinc  turnings 
by  treating  with  a  solution  of  mercuric  chloride,  and  then  add 
the  zinc  a  little  at  a  time  to  the  contents  of  the  flask.  On 
shaking  vigorously,  the  zinc  dissolves  completely  in  the  mercury 
and  then  becomes  oxidized  according  to  the  equation  given 
above.  To  show  the  presence  of  hydrogen  peroxide,  treat 
samples  of  the  resulting  turbid  liquid  with  an  acidified  iodide- 
starch  solution  containing  a  drop  of  ferrous  sulphate,  or  with  a 
titanium  sulphate  solution. 

(6)   Poly  sulphides. 
68.   Ammonium  Pentasulphide,  (NH4)2S5. 

The  sulphur  which  is  bound  directly  to  the  metal  in  metallic  sulphides  is 
capable  of  taking  on  more  sulphur  to  form  polysulphides  (Berzelius).  The 
additional  binding  power  of  the  sulphur  seems  to  be  influenced  to  a  high  degree 
by  the  metal  with  which  it  is  combined.  Polysulphides  of  the  alkali  metals 
are  numerous  and  well  known,  as  are  also  many  of  the  alkaline  earth  group; 
but  for  the  heavy  metals  either  none  exist  or  they  play,  at  most,  a  subor- 
dinate role.  Polysulphides  of  the  alkali  metals  can  be  obtained  by  fusion, 
(sulphur  livers),  those  of  caesium  and  rubidium  with  the  general  formula 
M2Sn,  in  which  n  =  2,  3,  4,  5,  and  6,  having  been  identified; '  polysulphides  can 
also  be  formed  by  dissolving  sulphur  in  solutions  of  the  monosulphides  and 
thereupon  crystallizing  or  precipitating  the  product.  (See  below.)  Of  the 
polysulphide  solutions,  those  of  the  tetrasulphides  with  the  ion  S4  are 
especially  stable.2  Besides  the  types  already  mentioned,  other  polysulphides 
have  been  obtained  under  varying  conditions  from  such  solutions,  but  the 
individuality  of  some  of  them  is  doubtful. 

Saturate  50  c.c.  of  concentrated  ammonia  solution  with  hydro- 
gen sulphide  in  a  closed  flask,  add  another  50  c.c.  of  the  ammonia, 
and  then  dissolve  in  this  mixture  as  much  roll  sulphur  as  pos- 
sible (about  50  g.)  at  30°  to  40°.  Filter  off  the  excess  of  sulphur 
and  to  the  yellow  solution  in  an  Erlenmeyer  flask  add  an  equal 
volume  of  95%  alcohol.  After  the  solution  has  stood  over  night 
in  the  ice-chest,  an  abundant  crystallization  of  intensely  yellow 
needles  of  ammonium  pentasulphide  is  obtained.  Drain  the  crys- 
tals, wash  them  with  alcohol  and  ether  and  allow  them  to  dry  for 
a  day  in  a  vacuum  desiccator  over  quicklime  upon  which  a  few 
drops  of  concentrated  ammonia  have  been  poured.  The  yield  is 

1  W.  Biltz  and  Wilke-Doerfurt,  Ber.  38,  123  (1905);  Z.  anorg.  Chem.  48, 
297;  50,  67  (1906). 

3  Kuster,  Z.  anorg.  Chem.  43,  53  (1904);  44,  431  (1905). 


AMMONIUM  TRIBROMIDE.  107 

about  40  g.  On  long  standing,  the  crystals  become  lighter  colored, 
due  to  decomposition.  Determinations  of  ammonia  and  sulphur 
confirm  the  above  formula. 

(c)    Polyhalides. 

Polysulphides  and  polyhalides  are  very  closely  related.  The  latter  likewise 
occur  chiefly  as  salts  of  the  alkali  and  alkaline  earth  metals  and  as  the  free 
halogen  acids.  The  stability  of  the  compounds  increases  very  rapidly  in  the 
series  of  the  alkali  metals  with  the  increase  in  the  atomic  weight  of  the  metal. 
From  caesium  and  rubidium  a  large  number  of  both  simple  and  mixed  polyha- 
lides of  the  type  MHa3  and  MHa5  (Ha  =  Cl,  Br,  or  I)  have  been  obtained.  As 
regards  the  stability  of  its  polyhalogen  compounds,  ammonium  comes  next  to 
the  higher  alkali  metals;  still  more  stable  are  the  substituted  ammonium  salts, 
namely,  the  polyhalides  of  the  complicated  alkaloids,  such  as  coniin,  nicotin, 
atropin,  narkotin,  and  of  the  diazonium  salts. 

69.    Ammonium  Tribromide,  NH4  [BrJ. 

Add  8  g.  of  bromine 1  to  a  lukewarm  solution  of  10  g.  ammonium 
bromide2  in  12  g.  of  water,  whereupon  the  temperature  of  the 
mixture  rises  a  little.  Allow  the  solution  to  stand  over  sul- 
phuric acid  in  a  vacuum  desiccator  which,  if  possible,  should  be 
placed  in  an  ice-chest.  Prismatic  crystals,  or  lamellar  aggre- 
gates, of  the  color  of  potassium  pyrochromate  separate  from  the 
solution  which  still  contains  free  bromine.  After  one  or  two  days 
drain  the  crystals  and  dry  them  in  a  vacuum  desiccator  con- 
taining a  small  dish  of  bromine  in  addition  to  sulphuric  acid  (the 
preparation  would  lose  bromine  to  an  atmosphere  free  from  that 
element).  The  yield  is  about  10  g.  Analyze  the  product  as 


1  This  is  only  half  the  calculated  amount.     The  use  of  more  bromine  does 
not  improve  the  yield. 

2  Ammonium  bromide  may  be  prepared  either  by  the  neutralization  of 
hydrobromic  acid  (No.  35)  with  ammonia  or  by  the  action  of  bromine  on 
ammonia.     According  to  the  latter  method,  allow  13  g.  bromine  to  flow  drop 
by  drop  from  a  dropp ing-funnel  into  130  c.c.  of  2-normal  ammonia  in  a  flask, 
which  is  surrounded  by  ice  and  is  constantly  shaken. 

8  NH3  +  3  Br2  =  6  NH4Br  +  N2. 

Evaporate  the  solution  on  the  water  bath,  and  if  desired  powder  the  residue 
and  dry  it  in  the  steam  closet.  The  salt  may  be  purified  by  recrystallization. 
Dissolve  it  in  its  own  weight  of  boiling  water,  cool  the  solution  with  ice,  and  add 
an  equal  volume  of  alcohol ;  recover  the  rest  of  the  salt  from  the  mother-liquor 
by  evaporating  and  crystallizing  in  a  similar  manner. 


108  LAW   OF  DISTRIBUTION. 

follows:  First  dissolve  a  sample  of  about  0.25  g.  in  a  solution 
of  potassium  iodide  and  titrate  with  0.1-normal  thiosulphate; 
second,  expose  about  0.25  g.  to  the  air  in  a  moderately  warm 
place  until  it  has  become  colorless,  and  weigh  the  residue  of 
ammonium  bromide  which  is  left. 

On  dissolving  the  tribromide  in  water,  free  bromine  separates 
in  considerable  quantity  and  one  molecule  of  Br2  can  be  com- 
pletely removed  by  shaking  the  solution  with  carbon  bisulphide, 
or  chloroform. 

70.   Law   of   Distribution;   Proof  of  the  Existence  of  Potassium 
Tribromide  and  Potassium  Tri-iodide. 

A.    Law  of  Distribution. 

If  to  two  non-miscible  solvents  standing  in  contact  with  one  another,  a  third 
substance  is  added  which  is  soluble  in  both,  the  ratio  of  the  concentration1 
of  this  substance  in  the  two  solvents  after  equilibrium  has  been  established, 
i.e.,  the  distribution  coefficient,  will  be  constant  for  a  given  temperature  pro- 
vided the  solute  has  the  same  molecular  weight  in  both  solutions: 

Cl  -k 
*' 

If  the  dissolved  substance  exists  in  one  of  the  two  solvents  in  a  disso- 
ciated or  associated  condition,  the  distribution  law  holds  only  for  the  same 
sort  of  molecular  aggregate  in  both  solvents.  This  fact  finds  expression  in 

c  n* 
the  formula  in  such  a  way  that  now  the  ratio  — —  remains  constant  if   nt 

C2  l 

and  n2  are  the  number  of  simple  molecules  associating  to  give  the  prevalent 
form  of  molecules  in  the  respective  solutions. 

I.  Dissolve  about  4  g,  of  succinic  acid  in  200  c.c.  of  water, 
shake  the  solution  with  50  c.c.  of  ether  in  a  500  c.c.  separatory 
funnel,  allow  the  two  liquids  to  separate,  and  titrate  10  c.c.  from 
both  layers  with  0.1-normal  sodium  hydroxide,  using  phenolphtha- 
lein  as  an  indicator.  To  what  is  left  in  the  separatory  funnel  add 
about  100  c.c.  of  water  and  some  ether,  shake  thoroughly  again, 
and  titrate  portions  of  20  c.c.  each  from  both  layers.  Again  add 
50  c.c.  of  water  and  half  as  much  ether,  and  determine  the  concen- 
trations in  20  c.c.  portions.  Finally  add  100  c.c.  of  water  and 
this  time  titrate  40  c.c.  portions.  The  four  distribution  coefficients 

1  Concentration  =  the  amount  of  substance  contained  in  a  unit  of  volume, 

m 


POTASSIUM    TRIBROMIDE.  109 

calculated  from  the  data  obtained  should  be  practically  equal. 
Succinic  acid  is  about  six  times  as  soluble  in  water  as  in  ether. 

II.  Dissolve  10  g.  of  benzoic  acid  in  a  mixture  of  100  g.  of 
water  and  100  g.  of  benzene;  shake  a  little  longer  after  it  is  all 
dissolved  until  the  distribution  equilibrium  is  reached,  and  titrate 
10  c.c.  from  each  layer.  Add  50  c.c.  of  water  and  again  shake 
vigorously,  repeating  this  several  times  in  succession  and  titrating 
both  layers  after  each  addition.  Nearly  constant  values  are 
obtained  in  this  experiment  if  the  solubility  in  water  is  divided 
each  time  by  the  square  root  of  the  solubility  in  benzene;  the 
molecular  weight  of  benzoic  acid  dissolved  in  benzene  is  nearly 
twice  as  large  as  in  water. 

B.  Proof  of  the  Existence  of  Potassium  Tribromide  in  Aqueous 
Solution. 

Bromine  dissolves  more  freely  in  a  potassium  bromide  solution  than  in 
water,  this  being  due  to  the  formation  of  the  potassium  salt  of  the  complex 
brom-hydrobromic  acid,  K  [BrBr^].  The  salt  has  never  been  obtained  in  a 
solid  form  suitable  for  analysis;  but  its  composition  can  be  derived,  with 
the  aid  of  the  mass-action  law  and  the  law  of  distribution,  by  finding  which 
of  the  assumptions,  x  =  2,  x  =  4,  etc.,  leads  to  conclusions  agreeing  with 
the  facts  found  experimentally.  Calculating  first  on  the  basis  of  the  com- 
plex ion  Br3~  (i.e.,  x  =  2),  this  ion  would  be  partially  broken  down  in  solution 
into  simple  bromine  ions  and  free  bromine: 

Br3~  <r>  Br~  +  Br2. 

For  the  state  of  equilibrium,  the  mass-action  law  gives: 

[Br-]  [Br2] 


The  three  concentrations  which  are  in  equilibrium  according  to  the  above 
equation  may  be  determined  by  preparing  a  concentrated  solution  of  bro- 
mine in  carbon  bisulphide  and  shaking  one  portion  of  it  with  water,  and  a 
second  portion  with  a  potassium  bromide  solution  of  known  molecular  con- 
centration,1 A.  In  each  experiment  the  molecular  quantities  of  bromine, 
D  and  B,  which  pass  into  the  aqueous  layer  are  determined;  the  excess  of 
bromine  in  the  second  case  is  equal  to  that  combined  in  the  complex:  thus, 
[Br3-]  =5  -  D. 

Furthermore,  [Br2]  =  D;  according  to  the  distribution  law,  the  amount 
of  free  bromine  existing  in  the  potassium,  bromide  solution,  as  well  as  that 


1  The  molecular  concentration  .is  the  amount  of  substance  dissolved  in 
one  liter  divided  by  its  molecular  weight. 


110  LAW  OF  DISTRIBUTION. 

existing  in  the  pure  water,  is  determined  solely  by  the  concentration  of 
bromine  in  the  carbon  bisulphide  layer. 

Finally,  [Br~]  =  A  —  (B  —  D),  i.e.,  the  equilibrium  concentration  of 
bromine  ions  is  equal  to  the  difference  between  the  original  concentration 
of  the  bromine  ions  and  that  of  the  part  which  forms  the  complex.  By 
inserting  these  values  in  the  mass-action  equation,  we  obtain: 

&D&-  (B-D)]  _ 

'B  -  D 

*i 

Mix  one  volume  of  bromine  with  four  volumes  of  carbon  bisul- 
phide and  add  5  c.c.  of  the  mixture  to  each  of  six  50  c.c.  glass- 
stoppered  bottles;  add  about  20  c.c.  of  water  to  each  of  two  of 
the  bottles,  and  to  the  other  four  add  equal  volumes  of  potassium 
bromide  solution  of  the  following  normal  concentrations: 

A  =  1/1;     1/2;     1/4;     1/8. 

Establish  equilibrium  between  the  layers  by  shaking  vigorously 
for  a  long  time;  place  the  bottles  in  water  at  room  temperature, 
and  give  them  an  occasional  lateral  motion  in  order  to  make  the 
drops  of  carbon  bisulphide  solution  floating  on  the  surface  sink. 
After  about  an  hour,  when  the  aqueous  layer  has  become  entirely 
clear,  pipette  off  10  c.c.  from  each  bottle,  and  after  adding  potas- 
sium iodide,  titrate  with  0.1-normal  thiosulphate  to  disappear- 
ance of  color.  In  this  way  the  concentrations  B  and  D  are 
obtained.  Substitute  these  values,  as  well  as  the  corresponding 
values  of  A,  in  the  mass-law  equation  and  see  whether  K  remains 
constant  throughout  the  experiment.  If  this  is  the  case,  it  proves 
the  existence  of  the  compound  KBr3.  It  is  a  good  plan  to  make 
the  assumption,  by  way  of  a  check,  that  more  than  one  molecule 
of  bromine  combines  with  the  potassium  bromide  in  the  formation 
of  the  complex,  to  substitute  the  values  in  the  corresponding 
mass-law  equation,  and  to  compare  the  values  of  K  thus  obtained. 
C.  Proof  of  the  Existence  of  Potassium  Tri-iodide  in  Aqueous 
Solution.  The  method  of  proof  is  the  same  here  as  in  the  above 
case.  Place  a  few  grams  of  finely  powdered  iodine  in  each  of 
seven  bottles,  and  add  250  c.c.  of  water  to  the  first,  250  c.c.  of 
T^s  normal  KI  to  the  second,  250  c.c.  of  ^  normal  KI  to  the 
third,  150  c.c.  of  &  normal  KI  to  the  fourth,  150  c.c.  of  &  normal 
KI  to  the  fifth,  75  c.c.  of  J  normal  KI  to  the  sixth,  and  75  c.c.  of 
\  normal  KI  to  the  seventh.  The  concentration  of  these  solutions 
must  be  accurately  known,  but  the  amount  taken  need  only  be 


SODIUM  HYDRAZOATE.  Ill 

roughly  measured.  Shake  the  seven  bottles  in  a  shaking-machine 
for  about  ten  hours  at  as  constant  a  temperature  as  possible.  Let 
settle,  and  titrate  100  c.c.  of  the  clear  solution  from  each  of  bottles 
1,  2,  and  3,  with  thiosulphate;  50  c.c.  each  from  bottles  4  and  5; 
25  c.c.  from  bottle  6;  and  20  c.c.  from  bottle  7.  From  these 
results  the  molecular  concentrations  of  iodine,  D  and  B,  are 
obtained;  the  value  of  A  is  the  original  concentration  of  the 
potassium  iodide  solution  used.  If  these  values  are  introduced 
in  the  above  mass-law  equation,  a  constant  value  for  K  should  be 
obtained,  which  at  20°  is  about  0.0013. 

Experiments  with  more  concentrated  potassium  iodide  solutions 
show  an  increase  in  the  values  of  the  constant;  this  indicates  the 
presence  of  higher  poly-iodides. 

71.  Rubidium  Iodide  Tetrachloride,  Rb  [IC1J;  Rubidium 

Tri-iodide,  RbI3. 

Dissolve  2.5  g.  of  rubidium  chloride  in  7.5  c.c.  of  water  and 
suspend  2.7  g.  of  iodine  in  the  solution.  On  passing  chlorine  into 
this  mixture  large,  beautiful,  orange-red  crystals  of  rubidium 
iodide  tetrachloride  are  obtained.  The  iodine  dissolves  during 
the  action  with  a  slight  liberation  of  heat,  and  later  the  new  salt 
separates  in  large  plates  which  increase  in  quantity  after  several 
hours  standing  in  the  ice-chest.  Drain  the  crystals  without  wash- 
ing and  allow  them  to  dry  for  an  hour  in  a  vacuum  desiccator  over 
sulphuric  acid.  Too  long  drying  causes  decomposition.  Yield 
about  5  g. 

Rubidium  Tri-iodide,  RbI3,  is  obtained  by  crystallization  from 
a  warm  solution  of  iodine  in  rubidium  iodide.  Neutralize  s 
solution  of  2.5  g.  of  rubidium  hydroxide  in  3  g.  of  water  with 
concentrated  hydriodic  acid,  add  6.2  g.  of  iodine,  and  heat  until 
solution  is  complete.  On  cooling,  the  salt  separates  in  crystals 
resembling  iodine.  Treat  them  as  in  the  preceding  preparation. 

(d)    Polynitrides. 

72.  Sodium  Hydrazoate  NaN3,  from  Sodamide  NaNH2. 

Sodamide  reacts  with  nitrous  oxide  to  form  sodium  hydrazoate,  the  salt 
of  hydrazoic  acid,  HN3;  from  an  aqueous  solution  of  sodium  hydrazoate, 
silver  nitrate  precipitates  the  difficultly  soluble  silver  salt,  a  substance  which 
like  the  pure  hydrazoic  acid  is  extremely  explosive.  Hydrazoic  acid,  or  azoi- 
mide,  was  discovered  by  Curtius  in  1890;  the  method  of  preparation  given 


112  POLYNITRIDES. 

below  was  devised  by  W.  Wislicenus,  in  1892.  The  relationship  between 
hydrazoic  acid,  iodo-hydriodic  and  brom-hydrobromic  acids  was  pointed  out 
by  Hantzsch  in  1895,  as  well  as  by  others. 

Sodamide.  Place  an  aluminium  boat  containing  about  3  g.  of 
sodium  (cf.  Sodium  Peroxide,  No.  65)  in  a  combustion  tube; 
support  the  tube  on  a  row  burner  and  heat  the  sodium  to  300-400° 
in  a  current  of  dry  ammonia.  Generate  this  gas  by  heating 
100  c.c.  of  concentrated  ammonia  solution  and  dry  it  by  passing 
through  a  tube  containing  soda-lime.  Hydrogen  escapes  from 
the  end  of  the  combustion  tube  together  with  the  excess  of 
ammonia,  and  the  mixture  can  be  made  to  burn.  Sodamide 
must  be  preserved  in  a  well-stoppered  bottle. 

Sodium  Hydrazoate.  Place  a  porcelain  boat  containing  0.5  g. 
sodamide  in  a  combustion  tube  which  is  inclosed  in  an  asbestos 
chamber  (see  p.  3).  Prepare  nitrous  oxide  by  heating  10  g.  of 
ammonium  nitrate  in  a  small  flask;  dry  the  gas  by  passing  it 
through  a  calcium  chloride  tube,  and  conduct  it  over  the  sodamide. 
Then  heat  the  combustion  tube  to  250°  until  the  reaction  is  finished, 
after  which  no  more  ammonia  can  be  detected  on  testing  the  escap- 
ing gases  with  moist  litmus  paper. 

NaNH2  +  N2O  =  NaN3  +  H2O, 
NaNH2  +  H20  =  NaOH  +  NH3. 

Silver  Hydrazoate.  Dissolve  the  sodium  hydrazoate  in  10  c.c.  of 
water;  a  portion  of  the  solution,  when  treated  with  a  little  ferric 
chloride,  gives  a  deep  brownish-red  color.  Acidify  another  and 
very  small  portion  of  the  solution  with  nitric  acid,  add  silver 
nitrate,  collect  the  precipitate  on  a  filter,  wash  it  with  water, 
alcohol  and  then  ether,  and  while  the  filter  is  still  moist  with  ether, 
tear  it  into  several  small  pieces.  After  drying  in  the  air  the 
small  amounts  of  the  preparation  adhering  to  the  bits  of  paper  will 
explode  violently  when  heated  or  struck.  It  is  very  dangerous  to 
prepare  any  larger  amount  of  silver  hydrazoate  than  that  indicated. 

Another  simple  method  for  obtaining  silver  hydrazoate  is  given 
by  Sabanejeff.  Heat  gently  1.5  g.  of  hydrazine  sulphate  (No.  122) 
with  4  c.c.  of  nitric  acid,  sp.  gr.  1.3,  in  a  test-tube  which  is  provided 
with  a  gas  exit  tube  bent  at  right  angles.  Pass  the  escaping  gas 
into  a  little  silver  nitrate  solution  in  a  second  test-tube;  silver 
hydrazoate  is  obtained  in  the  form  of  a  white,  curdy  precipitate. 
Filter  and  test  the  salt  as  before,  but  do  not  preserve  it. 


POTASSIUM  CYANATE.  113 

OXYACIDS   AND    THEIR   SALTS. 

(a)    Cyanates. 
73.   Potassium  Cyanate;  Urea  from  Ammonium  Cyanate. 

The  cyanates,  i.e.  the  salts  of  cyanic  acid,  are  produced  by  the  addition 
of  oxygen  to  the  cyanides: 

KCN  +  O  =  KCNO. 

With  regard  to  the  constitution  of  cyanic  acid  there  are  two  possible  for- 

mulas: 

x,N—  H  ^N 

I.        C^  II.        C 

^O  ^OH 

but  which  of  these  correctly  shows  the  structure  of  the  free  acid  and  of  its 
salts  has  not  been  definitely  determined;  it  is  quite  possible  that  both  forms 
exist  in  the  presence  of  one  another. 

Similarly  two  possible  formulas  may  be  written  for  hypochlorous  acid  : 


I.          C1  II.         Cl 

^0  \OH 

The  potassium  cyanide  used  in  the  preparation  of  the  cyanate  is  formed, 
together  with  a  little  cyanate,  by  heating  potassium  ferrocyanide  with  an 
alkali: 

K4[Fe(CN)6]  +  K2O  =  Fe  +  KCNO  +  5  KCN. 

In  the  presence  of  an  oxidizing  agent   (CrO3)   the  cyanide  is  oxidized  to 
cyanate. 

Dehydrate  130  g.  of  coarsely  broken  potassium  ferrocyanide  by 
stirring  it  in  a  shallow  iron  dish  over  a  rather  low  Fletcher  burner 
flame.  When  no  more  dark  yellow  particles  can  be  detected  in 
the  lumps,  grind  the  mass  to  a  fine  powder  and  remove  the  last 
traces  of  moisture  by  reheating.  While  still  warm,  triturate  100  g. 
of  the  powder  thus  obtained  with  75  g.  of  potassium  pyrochromate 
which  has  been  dried  by  being  melted.  Heat  the  mixture  in  the 
iron  dish  already  used,  whereupon  the  reaction  will  begin  in  spots 
and  spread  with  incandescence  throughout  the  mass.  Pulverize 
the  loose  black  product  while  still  warm,  cover  it  in  a  flask  with  a 
warm  mixture  of  450  c.c.  80  %  ethyl  alcohol  and  50  c.c.  of  methyl 
alcohol  and  boil  it  with  a  return  condenser  on  the  water  bath  for 
two  minutes.  Decant  the  hot  solution  through  a  plaited  filter  into 
a  beaker,  which  is  cooled  with  ice,  and  allow  the  salt  to  crystallize 
while  stirring.  Using  the  mother-liquor  from  this  crystallization 


114  OXY-HALOGEN  ACIDS. 

as  a  solvent,  extract  the  black  mass  again  in  the  same  manner  as 
before ;  filter,  and  repeat  the  extraction  three  or  four  times.  Collect 
all  of  the  cyanate  crystals  on  the  same  suction  filter,  and  after 
washing  with  ether,  dry  them  in  a  vacuum  desiccator  over  sul- 
phuric acid.  Yield,  30-40  grams.  The  product  may  be  used  in 
the  preparation  of  urea,  or  of  semicarbazid,  No.  123. 

Urea  from  Ammonium  Cyanate. 

Ammonium  cyanate  when  heated  in  aqueous  solution  undergoes  a  trans- 
formation into  urea  (Wohler,  1828). 

NH4CNO  =  CO(NH2)2. 

To  carry  out  this  classic  reaction,  evaporate  a  solution  of  8.1  g. 
potassium  cyanate  and  8.0  g.  ammonium  nitrate  to  dry  ness  on  the 
water  bath.  Boil  the  powdered  residue  in  a  flask  twice  with  alco- 
hol, and  concentrate  the  extract  until  a  crystallization  in  fine,  long 
needles  is  obtained.  Yield,  about  5  g. 

Heat  a  pinch  of  dry  urea  in  a  test-tube  until  it  just  melts,  and 
keep  it  at  this  temperature  for  about  a  minute;  ammonia  escapes. 
Dissolve  the  residue  in  a  little  water  and  add  a  drop  of  copper  sul- 
phate solution  and  some  sodium  hydroxide,  whereupon  a  rose- 
vioiet  coloration  appears;  this  is  the  so-called  biuret  reaction. 

(b)     Oxy-halogen  Acids. 

74.  Electrolytic  Production  of  Sodium  Hypochlorite  and  Potassium 

Chlorate. 

Sodium  hydroxide  and  chlorine  react  in  cold,  aqueous  solution,  forming 
sodium  hypochlorite,  sodium  chloride  and  water. 

2  NaOH  +  C12  =  NaCl  +  NaCIO  +  H2O. 

As  soon  as  chlorine  is  present  in  excess,  it  reacts  to  produce  free  hypochlo- 
rous  acid. 

HOH  +  C12  =  HOC1  +  HC1; 

and  the  latter  being  a  much  stronger  oxidizing  agent  than  sodium  hypochlo- 
rite, or  the  hypochlorite  ion,  it  oxidizes  chlorite  and  hypochlorite  ions  into 
chlorate  ions: 

2  HC1O  +  CIO-  =  ClO.r  -1-  2  HC1, 

3  HC1O  +  Cr  =  C1O3~  +  3  HC1. 

Heating  accelerates  these  reactions. 

The  raw  materials  necessary  for  these  experiments  can  be  prepared  more 
conveniently  by  the  electrolysis  of  alkali  chloride  solution  than  by  purely 


SODIUM    HYPOCHLORITE.  115 

chemical  means.  At  the  anode,  chlorine  is  the  primary  product;  sodium 
is  the  primary  product  at  the  cathode,  but  it  immediately  decomposes  water 
to  give  sodium  hydroxide  and  hydrogen.  The  products  formed  at  the  elec- 
trodes react  together  in  the  manner  shown  above  when  they  are  allowed 
to  mix  by  diffusion. 

The  process  as  outlined  above  is,  however,  interfered  with  somewhat  by 
the  progress  of  certain  secondary  reactions.  First,  the  hydrogen  produced 
at  the  cathode  reduces  the  hypochlorite,  and  to  some  extent  the  chlorate, 
to  chloride.  Since  only  the  discharged  hydrogen  atoms,  which  have  not 
yet  combined  to  form  molecules,  cause  the  reduction,  it  is  advantageous 
to  restrict  the  formation  of  hydrogen  to  as  small  an  area  as  possible;  by 
this  means  the  formation  and  escape  of  gaseous  hydrogen  is  favored.  In 
other  words,  the  current  density  at  the  cathode  must  be  kept  high.  To 
further  avoid  cathodic  reduction,  the  deposition  of  a  thin  skin  —  a  "dia- 
phragm "  —  of  hydrated  chromic  oxide  on  the  metal  of  the  cathode  works 
excellently;  this  can  be  most  simply  accomplished  by  the  electrolytic  reduc- 
tion of  a  little  alkali  chromate  which  is  added  to  the  electrolyte. 

Second,  the  C1O~  and  C1O3~  ions,  that  are  formed  in  the  process,  carry  a 
part  of  the  current,  and  when  they  become  discharged  at  the  anode,  they 
then  react  with  water  to  form  the  free  acids  and  oxygen.  The  current  which 
serves  to  discharge  these  ions  is,  therefore,  wasted.  The  loss  can  be  lessened 
by  using  a  high  anodic  current  density. 

Both  of  these  secondary  reactions  become  more  pronounced  as  the  elec- 
trolysis progresses,  i.e.,  as  the  concentration  of  the  chlorate  or  the  hypochlo- 
rate  becomes  greater.  This  explains  why  the  yield  for  a  given  amount 
of  current  gradually  grows  less  with  a  long-continued  electrolysis. 

Sodium  Hypochlorite.  A.  The  arrangement  of  the  electrical 
connections,  the  external  resistance,  and  the  measuring  instru- 
ments is  that  described  in  No.  14  (cf.  Fig.  9).  Place  the  beaker 
containing  the  electrolyte,  which  is  a  solution  of  88  g.  of  sodium 
chloride  in  500  c.c.  of  water  (3-normal),  inside  a  larger  beaker 
containing  ice  »water.  For  electrodes  use  two  sheets  of  platinum 
of  known  area,  e.g.,  30  sq.  cm.  The  current  density  should  be  the 
same  at  the  anode  as  at  the  cathode  and  about  15  amperes  per 
100  sq.  cm.  of  electrode  surface  (considering  only  one  side  of  the 
electrodes,  since  it  is  chiefly  between  the  inside  surfaces  that  the 
current  passes).  The  current  strength  holds  fairly  constant  and 
does  not  need  to  be  regulated  much  by  changing  the  resistance. 
The  minimum  potential  necessary  for  the  electrolysis  of  a  nor- 
mal sodium  chloride  solution  is  2.3  volts,  but  for  obtaining  the 
desired  current  density  at  least  6  volts  will  be  required.  The 
temperature  of  the  electrolyte  should  not  be  allowed  to  exceed 
20°.  Follow  the  extent  of  the  hypochlorite  formation  by  an 


116 


OXY-HALOGEN  ACIDS. 


analysis  every  10  to  20  minutes.  For  this  purpose,  remove  15  c.c. 
of  the  electrolyte  with  a  pipette,  let  it  stand  in  a  beaker  a  short 
time  until  the  gas  bubbles  have  escaped,  and  then  pipette  10  c.c. 
of  it  into  a  solution  of  potassium  iodide  which  is  slightly  acid  with 
hydrochloric  acid.  Titrate  the  iodine  set  free  with  0.1-normal 
thiosulphate  solution,  and  from  the  amount  required  calculate 
the  entire  amount  of  hypochlorite  present  in  the  whole  solution. 
To  determine  the  current  yield,1  find  the  total  number  of  ampere- 
seconds  used  at  the  time  of  taking  the  samples,  then  calculate 
from  this  the  theoretical  yield,  on  the  basis  that  96,540  amperes 
flowing  for  one  second  would  set  free  one  equivalent  each  of 
sodium  and  of  chlorine  and  thus  produce  one-half  mol.  of  NaClO. 
The  observed  data  and  the  calculated  values  can  be  arranged 
advantageously  in  a  table  as  follows: 


Time 
in  Min- 
utes. 

Volume 
in  c.c. 

Temper- 
ature. 

Amperes. 

Theoretical 
Yield  in 
Grams. 

c.c. 
Na2S203 
Used. 

Actual 
Yield  in 
Grams. 

Current 
Yield. 

0 

500 

10 

6. 

0 

0 

0 

0 

10 

500 

11 

6.6 

— 

— 

— 



20 

500 

13 

6.3 

2.91 

12.6 

2.34 

80.4% 

40 

485 

18 

6.2 

2.89 

19.8 

1.30 

45.0% 

Continue  the  electrolysis  until  the  current  yield  falls  below 
30%,  which  may  take  about  an  hour,  and  calculate,  from  the  last 
titration  showing  a  yield  better  than  30%,  the  number  of  grams 
of  "  active  "  chlorine  in  a  liter  of  the  solution  (7  to  9  g.). 

B.  Repeat  the  above  experiment,  using  again  500  c.c.  of  3-normal 
sodium  chloride  as  the  electrolyte,  but  adding  to  it  2.5  g.  of 
neutral  sodium  chromate.  Before  titrating  with  sodium  thio- 
sulphate, acidify  strongly  and  dilute  well,  in  order  that  the  color 
of  the  chromate  may  not  interfere  with  the  end-point.  From  the 
total  volume  of  thiosulphate  used  deduct  the  amount  which 
corresponds  to  the  iodine  set  free  by  the  chromate.  The  current 
yield  is  now  much  better,  on  account  of  the  addition  of  the  chro- 
mate, and  remains  above  30%  for  a  longer  time;  the  total  yield  of 
ft  active  "  chlorine,  computed  as  above,  now  amounts  to  from 
14  to  16  g.  per  liter. 

Compare  the  data  in  No.  14. 


POTASSIUM    CHLORATE.  117 

C.  If  experiment  A  is  repeated  without  cooling  the  solution, 
the  current  yield  falls  very  rapidly  to  below  30%,  and  the  quantity 
of  active  chlorine  produced  before  this  point  is  reached  amounts 
to  only  about  5  or  6  g.  per  liter. 

Test  qualitatively  the  bleaching  action  of  the  hypochlorite 
solutions  obtained  by  adding  a  little  to  some  indigo  solution. 

Potassium  Chlorate.  As  electrolyte  use  a  solution  containing 
100  g.  of  potassium  chloride  and  1  g.  of  potassium  pyrochromate 
in  250  c.c.  of  water.  Use  a  600  c.c.  beaker  for  the  electrolyzing 
vessel,  and  cover  it  with  the  two  halves  of  a  divided  watch  glass. 
The  electrodes  should  be,  as  before,  of  sheet  platinum.  Maintain 
at  the  anode  a  current  density  of  20  amperes  per  100  sq.  cm.  and  at 
the  cathode  a  higher  density.  Keep  the  temperature  at  40-60°, 
using  a  small  flame  if  necessary  to  add  to  the  heating  effect  of  the 
current.  A  source  of  current  at  from  6  to  10  volts  will  suffice  to 
maintain  the  necessary  potential.  Pass  a  slow  stream  of  carbon 
dioxide  continuously  into  the  solution  between  the  electrodes.  To 
obtain  a  good  yield,  about  60  ampere-hours  are  required;  thus 
for  an  anode  surface  (one  side)  of  25  sq.  cm.  the  experiment  must 
be  continued  14  hours.  If  necessary  it  is  permissible  to  inter- 
rupt the  electrolysis.  Occasionally  replace  the  water  lost  by 
evaporation. 

A  little  potassium  chlorate  crystallizes  out  during  the  electroty- 
sis,  but  the  main  part  is  obtained  after  cooling;  drain  the  crystals 
with  suction  and  wash  with  a  little  cold  water.  To  determine  the 
entire  yield  of  chlorate,  dilute  the  mother-liquor  to  500  c.c.  and 
titrate  a  part  of  it  with  ferrous  sulphate:  boil  10  c.c.  of  the  solution 
in  a  flask  in  order  to  drive  out  the  free  chlorine,  replace  the  air  by 
means  of  carbon  dioxide,  and  after  cooling  add  50-70  c.c.  of  a 
0.1-normal  ferrous  ammonium  sulphate  solution  acidified  with 
sulphuric  acid.  Close  the  flask  with  a  Bunsen  valve  and  boil  the 
solution  ten  minutes.  After  cooling,  dilute  the  liquid  to  twice  its 
volume,  add  20  c.c.  of  a  20%  manganous  sulphate  solution  to  pre- 
vent the  hydrochloric  acid  from  interfering  with  the  titration,  and 
titrate  in  the  cold  with  0.1-normal  potassium  permanganate.  A 
slight  correction  can  be  applied  to  allow  for  the  chromate  present. 
Determine  in  a  similar  manner  the  percentage  of  KC1O3  in  the 
crystals  obtained;  then  calculate  the  entire  quantity  of  potassium 
chlorate  and  the  current  yield.  It  is  to  be  remembered  that 


118  OXY-HALOGEN  COMPOUNDS. 

according  to  the  equation  for  its  formation,  96,540  ampere-seconds 
yield  but  1/6  of  a  mol.  of  KC103.  The  current  yield  amounts  to 
about  70%,  and  about  85%  of  all  the  chlorate  is  obtained  in  the 
crystals.  A  further  loss  occurs  when  the  crude,  solid  product  is 
purified  by  recrystallization  from  hot  water.  The  yield  of  puri- 
fied chlorate  actually  obtained  in  an  experiment  carried  out 
according  to  the  above  directions,  was  23  g. 

The  chlorate  dissociates  into  a  potassium  ion  and  a  chlorate  ion. 
Test  the  purity  of  the  preparation  by  dissolving  a  little  in  water, 
acidifying  with  nitric  acid  and  adding  silver  nitrate;  there  should 
be  no  precipitate  of  silver  chloride. 

To  illustrate  the  decomposition  of  chlorates  by  heat,  melt  a  little 
of  the  preparation  in  a  test-tube  and  test  for  oxygen  with  a  glowing 
splinter  (see  the  next  preparation). 

75.   Potassium  Perchlorate. 

Potassium  chlorate  on  being  heated  to  about  400°  decomposes  in  two 
different  ways: 

1.  4  KC1O3  =  KC1  +  3  KC1O4. 

2.  KC103  =  KC1  +  3  O. 

If  the  vessel  is  clean  and  the  potassium  chlorate  pure,  the  decomposition 
proceeds  essentially  according  to  equation  1.  If,  on  the  other  hand,  cataly- 
zers are  present,  such  as  manganese  dioxide  or  ferric  oxide,  or  if  it  is  heated 
to  a  higher  temperature,  then  the  reaction  takes  place  principally  according 
to  equation  2.  It  is  on  this  account  that  a  mixture  of  manganese  dioxide 
with  potassium  chlorate  is  used,  rather  than  the  pure  chlorate,  in  pre- 
paring oxygen. 

Heat  50  g.  of  potassium  chlorate  in  a  new  100  c.c.  porcelain  cru- 
cible until  the  salt  just  melts.  Without  increasing  the  heat  keep 
the  temperature  as  uniform  as  possible,  so  that  oxygen  barely 
escapes  while  the  melt  gradually  becomes  more  viscous  and  pasty. 
If  at  the  end  of  10  or  15  minutes  the  mass  has  become  uniformly 
semi-solid,  allow  it  to  cool,  then  cover  it  with  50  c.c.  of  cold  water 
and  allow  it  to  stand  until  fully  disintegrated.  Collect  the  undis- 
solved  potassium  perchlorate  on  a  filter  and  recrystallize  it  from 
200  c.c.  of  water.  Yield,  30  g.  The  potassium  chloride  passes  into 
the  filtrates.1 


1  At  the  room  temperature  about  36  g.  KC1,  6.6  g.  KC1O3,  and  1.5  g. 
KC1O4  are  soluble  in  100  c.c.  of  water. 


IODIC  ACID.  119 

Potassium  perchlorate,  like  the  chlorate,  yields  no  precipitate 
with  silver  ions. 

Heat  in  a  dry  test-tube  a  mixture  of  potassium  chlorate  and  one- 
sixth  of  its  weight  of  manganese  dioxide,  and  test  the  oxygen 
evolved  with  a  glowing  splinter.  Observe  that  the  evolution  of 
oxygen  takes  place  at  a  lower  temperature  than  when  pure  potas- 
sium chlorate  is  used.  Extract  the  residue  with  water,  and  filter; 
potassium  chloride  is  in  the  filtrate,  as  may  be  shown  with  silver 
nitrate. 

Free  perchloric  acid  is  not  explosive.  Potassium  perchlorate 
when  covered  with  concentrated  sulphuric  acid  remains  unchanged 
in  the  cold  and  does  not  explode  on  gentle  heating;  potassium 
chlorate,  on  the  other  hand,  yields  explosive  chlorine  dioxide  even 
in  the  cold. 

76.    lodic  Acid  and  lodic  Anhydride ;  a  "  Time  Reaction." 

Seal  a  0.5-1.0  meter  long  glass  extension  tube  to  the  neck  of  a 
round-bottomed  flask,  and  boil  32  g.  of  iodine  with  130  g.  of  con- 
centrated, colorless  nitric  acid  in  the  flask,  using  a  ring  burner  to 
avoid  bumping.  Remove  the  lower  oxides  of  nitrogen,  as  fast  as 
they  are  formed,  by  means  of  a  current  of  carbon  dioxide,  or  of  air. 
After  the  oxidation  is  complete  and  the  solution  has  cooled,  collect 
the  solid  iodic  acid  on  an  asbestos  filter  and  separate  it  from  the 
asbestos  fibers  by  dissolving  in  the  least  quantity  possible  of  hot 
water  and  filtering.  Allow  the  iodic  acid,  HIO3,  to  crystallize, 
after  concentrating  somewhat  if  necessary,  by  letting  the  solution 
stand  in  a  vacuum  desiccator  over  sulphuric  acid.  Evaporate  the 
mother-liquor  and  dehydrate  the  residue  at  200°.  Iodic  anhy- 
dride, I2O5,  is  thus  obtained. 

Heat  a  sample  of  the  iodine  pentoxide  in  a  test-tube;  it  is  broken 
down  into  iodine  and  oxygen,  as  is  shown  by  the  violet  vapors  and 
by  the  test  with  a  glowing  splinter. 

"  Time  Reaction." 

If  a  solution  of  iodic  acid  is  added  drop  by  drop  to  an  aqueous  solution  of 
sodium  sulphite  acidified  with  sulphuric  acid,  the  solution  at  first  remains 
vx)lorless: 

3  H2SO3  +  HIO3  =  3  H2S04  +  HI. 


120  OXY-HALOGEN    COMPOUNDS. 

When,  with  further  addition  of  iodic  acid,  all  the  sulphurous  acid  becomes  oxi- 
dized, the  reduction  then  continues  at  the  expense  of  the  hydriodic  acid 
which  has  been  formed : 

HIO3  +  5  HI  =  3  H2O  +  3  I2. 

Thus  if  an  acidified  sulphite  solution  is  treated  with  more  than  one-third  of  a 
mol.  of  iodic  acid,  a  separation  of  iodine  occurs;  if  the  solution  is  concen- 
trated this  takes  place  immediately;  if  dilute,  only  after  some  time,  and  then 
suddenly  and  completely. 

This  very  remarkable  retardation  reminds  one  of  the  phenomena  of  super- 
cooling, superheating,  and  supersaturation,  but  it  is  different  inasmuch  as  the 
reaction  takes  place  in  a  homogeneous  medium,  and  is  not  followed  by  a 
separation  of  material  in  another  state  of  aggregation. 

Prepare  one  solution  by  dissolving  1.8  g.  of  iodic  acid,  or  an 
equivalent  amount  of  potassium  iodate,  in  water  and  diluting  to 
one  liter;  prepare  a  second  solution  with  0.9  g.  of  Na2S03.7  H20, 
5  g.  of  10%  sulphuric  acid,  and  9.5  g.  of  starch  (the  latter  should 
first  be  suspended  in  a  little  cold  water  and  then  stirred  into  a 
beaker  of  boiling  water  to  form  a  paste),  and  dilute  this  likewise 
to  one  liter.  Measure  100  c.c.  of  each  solution  into  separate 
beakers,  mix  these  two  portions  at  a  definite  instant,  and  count 
the  seconds  until  a  deep  blue  color  suddenly  appears.  Check 
this  result  by  repeating  the  experiment.  Then  dilute  each  of  the 
solutions  to  f ,  f ,  f ,  and  £  of  their  former  concentrations,  and  deter- 
mine for  each  dilution  the  time  which  elapses  before  the  blue  color 
appears.  Plot  graphically  the  dependence  of  the  time  on  the  dilu- 
tion. The  results  are  reproducible  and  comparable  only  when  the 
temperature  of  the  solutions  does  not  vary  appreciably. 

77.   Potassium  Iodate  from  Potassium  Chlorate. 

KC103+  I  =  KI03+  Cl. 

Place  30  g.  of  potassium  chlorate  in  a  200  c.c.  flask  and  dissolve 
it  in  60  c.c.  of  warm  water.  Add  35  g.  of  iodine  to  the  solution, 
and  while  maintaining  the  mixture  at  a  moderate  temperature 
introduce  1  to  2  c.c.  of  concentrated  nitric  acid.  (Hood.)  After 
one  or  two  minutes  a  vigorous  reaction  begins  and  a  stream  of 
chlorine  escapes,  carrying  with  it  a  little  iodine.  When  the  reaction 
moderates,  boil  to  drive  off  the  chlorine,  and,  when  this  is  nearly 
accomplished,  add  1  g.  more  of  iodine.  Concentrate  by  evapora- 
tion until  on  cooling  nearly  all  of  the  potassium  iodate  crystallizes; 


SODIUM  NITRITE.  121 

collect  the  product  on  a  filter,  and  recover  what  is  left  in  the 
mother-liquor  by  evaporation. 

Dissolve  the  crude  product,  which  always  contains  some  acid 
salt,  in  150  c.c.  of  hot  water,  and  neutralize  exactly  with  potassium 
hydroxide.  On  cooling  a  good  yield  of  the  pure  salt  is  obtained. 

Ignite  a  little  of  the  product  in  a  porcelain  crucible,  and  test  the 
residue  for  chloride  by  distilling  it  with  potassium  bichromate  and 
concentrated  sulphuric  acid,  and  passing  the  vapors  into  ammonia 
water. 

78.   Potassium  Bromate  and  Potassium  Bromide. 

To  a  solution  of  62  g.  potassium  hydroxide  in  62  g.  of  water, 
add  80  g.  of  bromine  drop  by  drop  while  cooling  by  means  of  tap 
water.  (Hood.)  The  solution  soon  becomes  colored  a  permanent 
yellow,  and  later  a  crystalline  powder  of  potassium  bromate  sepa- 
rates; after  cooling  completely,  collect  the  bromate  on  a  filter  and 
purify  it  by  recrystallization  from  130  c.c.  of  boiling  water.  Com- 
bine all  the  mother -liquors  and  evaporate  them  to  a  semi-solid 
mass;  mix  this  thoroughly  with  5  g.  of  powdered  wood-charcoal, 
dry  it  completely  and  then  heat  it  to  redness  for  an  hour  in  a  large 
porcelain  crucible  surrounded  by  an  asbestos  or  sheet-iron  funnel. 
Treat  the  sintered  mass  with  120  c.c.  of  hot  water,  and  then 
wash  the  residue  with  20  c.c.  more  of  water;  evaporate  the  filtered 
solution  to  crystallization.  The  yield  is  26  to  27  g.  of  KBrO3 
and  90  to  95  g.  of  KBr. 

The  addition  of  acid  to  the  aqueous  solution  of  either  one  of 
these  salts  should  not  produce  a  yellow  coloration,  due  to  the 
separation  of  free  bromine. 

(c)    Nitrites  and  Nitrates. 
79.    Sodium  Nitrite  from  Sodium  Nitrate. 

Sodium  nitrate,  when  melted  with  a  reducing  agent  such  as  lead,  loses 
one-third  of  its  oxygen  and  goes  over  into  the  nitrite. 

In  the  absence  of  reducing  agents  sodium  nitrate  can  be  melted  without 
decomposition;    at  higher  temperatures,   however,  a  dissociation,  although 
incomplete,  takes  place  according  to  the  equation: 
NaN03  =  NaN02  +  O. 

From  the  above  facts  the  conclusion  may  be  drawn  that  this  dissociation  of 
sodium  nitrate  takes  place  even  at  more  moderate  temperatures,  although 
to  so  small  extent  that  it  can  be  proved  only  indirectly.  The  reducing  agent 
by  removing  oxygen,  one  of  the  products  of  the  dissociation,  causes  the 


122 


NITRITES  AND  NITRATES. 


decomposition  to  continue  until  it  has  become  of  appreciable  magnitude. 
This  is  a  good  example  of  a  non-electrolytic  dissociation  which  though 
actually  insignificant  can  be  made  apparent  by  the  use  of  a  reagent. 

Sodium  nitrite  is  the  most  important  technically  of  all  the  salts  of  nitrous 
acid.  It  is  used  principally  for  diazotizing  in  the  manufacture  of  azo-dyes. 

Heat  85  g.  of  sodium  nitrate  in  an  iron  dish,  of  15  cm.  diameter, 
until  it  melts,  and  add  207  g.  of  lead  a  little  at  a  time  while  stirring 
with  an  iron  spatula.  Continue  the  heating  until  all  the  lead  is 
oxidized,  which  may  take  half  an  hour,  and  then  while  cooling 
keep  stirring  in  order  to  obtain  the  mass  in  small  loose  lumps. 
Extract  the  product  first  with  300  c.c.  of  hot  water  and  then  twice 
more  with  100  c.c.  To  precipitate  the  lead  which  has  gone  into 
the  solution  as  plumbite,  pass  in  carbon  dioxide  for  a  few  minutes. 
Filter  and  neutralize  the  filtrate  cautiously  with  a  very  little  dilute 
nitric  acid.  Evaporate  to  a  small  volume  to  obtain  crystals,  collect 
these  on  a  suction-filter,  wash  with  alcohol,  and  evaporate  the 
mother-liquor  to  obtain  more  crystals.  Yield,  40-50  g. 

The  lead  oxide  can  be  again  converted  into  lead  by  reduction 
with  charcoal  (No.  1). 

80.   Potassium  Nitrate  from  Sodium  Nitrate. 

If  a  mixture  of  sodium  nitrate  and  potassium  chloride  is  boiled  with  a 
quantity  of  water  insufficient  for  complete  solution,  the  undissolved  residue 
will  be  sodium  chloride,  which  is  the  least  soluble  in  hot  water  of  the  four 
possible  salts: 

NaNO3  +  KC1  +±  KNO3  +  NaCl. 

If  the  filtrate  is  cooled,  potassium  nitrate  will  crystallize,  since  this  is  the 
least  soluble  at  the  room  temperature. 

GRAMS  OF  SALT  SOLUBLE  IN   100  GRAMS  OF  WATER. 


At  10°. 

At  100°. 

Potassium  nitrate  
Sodium  chloride  
Potassium  chloride.  ..  .  . 
Sodium  nitrate 

21 
36 
31 

81 

246 
40 
56 
180 

Thus  the  interaction  of  the  salts  used  depends  entirely  upon  the  solu- 
bility relations. 

This  method  has  had  an  important  application  in  the  manufacture  of 
potassium  nitrate  for  the  gunpowder  industry. 

Dissolve  190  g.  of  crude  Chile  saltpeter  in  200  c.c.  of  boiling 
water  in  a  previously  weighed  flask.  To  the  boiling  solution 


SILVER   NITRATE.  123 

add  150  g.  of  powdered  potassium  chloride  and  boil  for  half  an 
hour  longer,  replacing  any  water  lost  by  evaporation.  The  con- 
tents of  the  flask  should  at  the  end  weigh  520  to  540  g.  While  still 
hot,  filter  rapidly  through  a  Biichner  funnel  and  rinse  the  residue 
with  a  test  tube  full  of  hot  water.  Cool  the  filtrate  rapidly  while 
shaking,4  whereby  a  crystalline  meal  of  potassium  nitrate  is  formed. 
Evaporate  the  mother-liquor;  remove  the  sodium  chloride  by 
filtering  while  hot;  and  cool  the  filtrate  rapidly  to  obtain  more 
potassium  nitrate.  If  sufficient  mother-liquor  still  remains,  work 
it  up  in  the  same  manner  to  obtain  a  further  yield.  Unite  all  of 
the  crystals  of  potassium  nitrate  and  purify  them  by  recrystalliza- 
tion  until  they  are  free  from  chloride.  The  yield  is  from  60  to  70% 
of  the  calculated. 

81.    Silver  Nitrate. 

The  nitrates  of  the  alkali  metals  break  down  into  nitrites  when  heated; 
those  of  all  the  other  metals  dissociate  into  metallic  oxide  and  nitric  anhy- 
dride or  its  decomposition  products: 

Cu(NO3)2  =  CuO  +  N2O4  +  O. 

This  dissociation  begins  to  take  place  at  very  different  temperatures  with 
the  various  metals.  Silver  nitrate  can  be  melted  without  decomposition, 
while,  at  the  same  temperature,  the  nitrates  of  metals  with  a  higher  valence, 
for  example  copper  nitrate,  are  decomposed;  thus  by  melting  a  mixture  of 
these  two  nitrates  and  dissolving  the  fusion,  silver  nitrate  can  be  obtained 
free  from  copper. 

By  carefully  regulating  the  temperature  and  repeating  the  process,  mix- 
tures of  very  closely  related  nitrates  can  be  separated;  the  "nitrate  method" 
for  separating  the  metals  of  the  rare-earths  depends  upon  this  principle. 

A  completely  analogous  behavior  is  shown  by  the  sulphates,  as  in  the 
Ziervogel  process  for  obtaining  silver  from  argentiferous  pyrite.  By  roasting, 
the  sulphates  of  the  metals  are  first  formed,  all  of  which,  however,  with  the 
exception  of  silver  sulphate,  decompose  at  a  somewhat  higher  temperature 
into  metal  oxide  and  sulphur  trioxide;  by  leaching,  the  silver  is  obtained  in 
the  solution. 

Dissolve  a  silver  coin  in  30%  nitric  acid,  evaporate  the  solution 
to  dry  ness,  and  transfer  the  residue  to  a  porcelain  crucible.  Place 
this  crucible  upon  a  wire  triangle  inside  a  larger  crucible,  and 
gradually  heat  the  outer  crucible  to  a  dull  red  heat.  When  the 
decomposition  is  completed,  extract  the  black  residue  with  water, 
concentrate  the  filtrate  and  test  it  for  copper.  Should  any  copper 
be  present,  which  will  frequently  be  the  case,  evaporate  the  solu- 


124  MANGANATES  AND  FERRATES. 

tion  to  dry  ness  and  heat  the  silver  nitrate  carefully  until  it  just 
begins  to  melt.  In  this  way  the  remainder  of  the  copper  nitrate 
is  decomposed.  Repeat  the  extraction  with  water. 

The  copper  oxide  residues  should  contain  very  little  silver. 

82.   Bismuth  Nitrate  and  Basic  Bismuth  Nitrate. 

Bismuth  nitrate  may  be  obtained  from  the  solution  of  the  metal  in  nitric 
acid  in  the  form  of  large,  colorless  crystals  with  5  molecules  of  water.  The 
Bait  is  strongly  hydrolyzed  by  water,  and,  according  to  the  temperature  and 
the  concentration  of  the  acid,  basic  salts  of  various  compositions  are  pro- 
duced.1 The  following  directions  yield  a  precipitate  of  approximately  the 
composition,  4  Bi2O3 .  3  N2O5 .  9  H2O,  which  is  not,  however,  to  be  regarded  as 
a  homogeneous  compound. 

Dissolve  100  g.  of  coarsely  pulverized  bismuth  by  heating  it  in 
a  flask  with  500  g.  of  nitric  acid  (sp.  gr.  1.2).  Filter  through  a 
hardened  filter,  using  suction,  and  evaporate  in  a  porcelain  dish 
until  crystallization  begins.  Collect  the  crystals  on  a  suction- 
filter,  wash  with  a  little  nitric  acid  (sp.  gr.  1.2)  and  dry  in  a  desic- 
cator. Evaporate  the  mother-liquor  to  obtain  more  crystals. 

To  prepare  the  above  mentioned  basic  salt,  triturate  one  part  of 
the  bismuth  nitrate  with  four  parts  of  water  and  stir  this  mixture 
into  21  parts  of  boiling  water.  Allow  the  precipitate  to  settle, 
wash  by  decantation,  collect  on  a  suction-filter  and  dry  the  prepa- 
ration at  a  temperature  not  exceeding  30°. 

(d)    Manganates  and  Ferrates. 
83.   Potassium  Permanganate  by  the  Fusion  Method. 

By  fusing  manganese  compounds  in  an  oxidizing-alkaline  flux  a  man- 
ganate  is  formed,  and  this,  when  it  is  dissolved  in  water  and  the  free  alkali 
is  neutralized,  changes  into  permanganate  and  manganese  dioxide: 

MnO2  +  K2CO3  +  O  =  CO2  +  K2MnO4, 
3  K2MnO4  +  2  H2O  =  MnO2  +  4  KOH  +  2  KMnO4. 

Melt  together  80  g.  of  potassium  hydroxide  and  40  g.  of  potas- 
sium chlorate  in  a  sheet  iron  crucible  6  to  8  cm.  in  diameter. 
Remove  the  flame  and  while  stirring  with  an  iron  spatula  (a  heavy 
wire  or  an  old  file),  add  80  g.  of  finely  powdered  pyrolusite,  quite 
rapidly  but  not  all  at  once.  The  fusion  effervesces  somewhat. 
Heat  again,  at  first  moderately  then  more  strongly,  and  stir 

1  Cf.  A.  Findlay:  The  Phase  Rule. 


POTASSIUM  PERMANGANATE.  125 

vigorously  all  the  while  until  the  mass  has  become  dry.  Finally 
heat  for  5  minutes  at  a  dull  red  heat.  Unless  stirred  as  directed, 
the  melt  will  solidify  to  a  hard  cake  which  can  be  removed  from 
the  crucible  only  with  difficulty.  When  it  is  cold,  break  up  the 
mass  and  boil  it  with  1.5  liters  of  water  while  conducting  a  vigorous 
stream  of  carbon  dioxide  into  the  liquid.  When  the  manganate 
is  completely  decomposed  and  a  drop  of  the  solution  gives  a  clear, 
violet-red  spot  on  filter  paper,  with  no  trace  of  green,  allow  the 
precipitate  to  settle,  and  decant  the  liquid  as  carefully  as  possible 
from  the  sludge  of  manganese  dioxide.  Filter  with  suction 
through  a  felt  of  asbestos  on  a  Biichner  funnel.  Concentrate  the 
filtrate  to  one-half,  filter  again  through  asbestos,  and  evaporate 
until  crystallization  begins.  Collect  the  crystals  on  a  porcelain 
filter-plate  and  wash  with  a  little  cold  water.  Obtain  a  second 
crop  of  crystals  from  the  mother  liquor  and  examine  with  a  micro- 
scope to  see  if  it  is  free  from  crystals  of  potassium  chloride. 
Finally  recrystallize.  The  yield  should  be  50  to  60  g.  of  the  pure 
salt. 

84.  Electrolytic  Preparation  of  Potassium  Permanganate. 

If  a  solution  of  an  alkali  hydroxide,  or  better  of  an  alkali  carbonate,  is 
electrolyzed  with  an  anode  of  manganese,  permanganate  is  formed  by  anodic 
oxidation.  For  the  reaction  to  succeed  it  must  be  carried  out  at  a  rather 
high  temperature,  since  otherwise  only  manganate  is  formed.  For  the  cathode, 
nickel  or  iron  wires  may  be  used,  but  they  must  be  surrounded  with  a  porous 
clay  diaphragm  to  avoid  reduction  of  the  permanganate. 

Place  a  porous  cup,  10  cm.  high  and  4  cm.  wide,  in  a  liter 
beaker,  and  inside  the  cup  place  an  iron  wire  to  serve  as  the 
cathode.  For  the  anode  use  a  good  sized  lump  of  manganese 
(obtained  by  the  Goldschmidt  process;  cf.  No.  2);  fasten  it  by 
means  of  a  fine  platinum  wire  to  a  stout  iron  wire  and  allow  only 
the  manganese  to  dip  into  the  liquid.  Use  as  the  electrolyte 
200  c.c.  of  a  solution  of  potassium  carbonate  saturated  at  0° 
(solubility:  105  g.  of  K2CO3  in  100  g.  of  water)  which  should  stand 
at  the  same  level  inside  and  outside  the  porous  cell.  In  order  to 
effect  a  rapid  electrolysis  it  will  probably  be  necessary,  on  account 
of  the  high  resistance  of  the  diaphragm,  to  draw  from  a  current 
supply  of  more  than  6  volts.  Raise  the  temperature  of  the  bath 
to  50-60°  and  maintain  it  at  this  point  by  using  an  external  flame 
whenever  it  is  necessary  to  supplement  the  heating  effect  of  the  cur- 


126  OXY-ACIDS   OF  SULPHUR. 

rent.  The  formation  of  permanganate  begins  as  soon  as  the  cur- 
rent starts ;  red  films  of  solution  form  at  the  anode  surface  and  sink 
to  the  bottom  of  the  beaker.  At  the  end  of  30  minutes,  and  then 
after  every  20  minutes,  stir  the  bath  well  and  take  samples  of  5  or 
10  c.c.  with  a  pipette  for  the  purpose  of  determining  the  current 
yield.1  Acidify  the  sample  with  sulphuric  acid,  whereupon  the 
precipitate,  which  is  mainly  ferric  hydroxide,  dissolves;  decolorize 
the  hot  solution  with  a  measured  excess  of  0.1-normal  oxalic  acid  * 
and  titrate  back  at  60°  with  0.1-normal  permanganate.  The  per- 
manganate solution  obtained  in  this  manner  is  so  dilute  that  it 
does  not  pay  to  attempt  to  crystallize  the  solid  salt.  The  current 
yield  in  several  experiments  made  in  the  author's  laboratory  with 
1.5-3.0  amperes  was  as  high  as  25%  at  the  end  of  45  minutes;  in 
the  next  hour  it  sank  to  about  10%. 

85.   Barium  Ferrate,  Ba[FeOj. 

Clamp  a  50-75  c.c.  flask  in  an  upright  position  and  place  in  it  a 
mixture  of  10  g.  of  fine  iron  filings  and  20  g.  of  potassium  nitrate. 
Heat  with  a  flame  until  a  reaction  takes  place  and  the  iron  burns 
with  a  shower  of  sparks.  After  cooling  break  the  flask,  extract 
four  times  with  50  c.c.  of  ice  water,  and  filter  the  combined  extract 
through  asbestos.  Pass  the  deep  violet-red  filtrate  again  through 
a  fresh  asbestos  felt  and  precipitate  it  immediately  with  a  solution 
of  barium  chloride.  After  an  hour  drain  the  precipitate  of  barium 
ferrate  on  a  hardened  filter,  wash  it  with  alcohol,  then  with  ether, 
and  dry  at  the  room  temperature  in  a  vacuum  desiccator.  Yield, 
1.0  to  1.5  g. 

(e)    Oxy-acids  of  Sulphur  and  Their  Salts. 

For  sulphurous  acid  the  two  following  constitutional  formulas2  have  been 
proposed : 

O^     /OH  /OH 

0^S\H  *\OH 

Unsymmetrical  Formula.  Symmetrical  Formula. 


1  Compare  the  directions  in  Preparations  14  and  74. 

2  Constitutional  formulas  of  the  complex  inorganic  acids  apply  only  with 
a  certain  reservation:   They  characterize  the  ability  to  react  only  in  one  par- 
ticular direction,  and   show,  of  several  possibilities,  only  one  definite   con- 
dition of  the  acid  (see  cyanic  and  hypochlorous  acids).     They  are  of  value, 
however,  as  aids  to  the  memory. 


OXY  -ACIDS  OF  SULPHUR.  127 

For  a  further  discussion  of  these  formulas,  see  Nos.  155  and  156. 
Regarding  the  constitution  of  sulphuric  acid,  it  is  to  be  noted: 

1.  Sulphuric  acid   contains  two  hydroxyl   radicals:     sulphuryl    chloride 
(No.   149)  and  water  react  together  with  the  formation  of  sulphuric  acid, 
whereby  the  chlorine  atoms  are  replaced  by  hydroxyl  groups: 

SO2C12  -»  SO2(OH)C1  ->  SO2(OH)2. 

2.  The  two  hydroxyl  groups  of  sulphuric  acid  are  bound  to  the  sulphur 
atom:    Either  one  or  both  of  these  groups  can  be  replaced  by  organic  radi- 
cals, such  as  phenyl,  whereby  phenyl  sulphonic  acid,  C6H5SO2OH,  and  diphe- 
nylsulphone,  (C6H5)2SO2,  are  formed.     But  the  same  substances  can  be  pre- 
pared by  the  oxidation  respectively  of  mercaptan,  C6H5SH,  and  of  diphenyl 
sulphide,  (C6HS)^,  in  which  the  organic  radical  must  be  joined  directly  to 
the  sulphur  atom.     From  this  it  follows  that   both  hydroxyl  groups  are 
held  by  the  same  sort  of  a  bond  : 

/OH 

o2s; 

\OH 

3.  The  way  in  which  the  two  remaining  oxygen  atoms  are  bound  is  not 
known  writh  certainty,  but  a  mutual  bonding  of  the  two  atoms  to  one  another 
is  improbable  on  account  of  the  great  stability  of  the  SO2-grouping.     Thus 
the  structural  formula  is  apparently 


PYROSULPHURIC  ACID,  HO.SO2-O-SO2.OH,  may  be  regarded  as  derived 
from  two  molecules  of  sulphuric  acid  by  the  loss  of  one  molecule  of 
water;  it  is  prepared  by  dissolving  sulphur  tri  oxide  in  sulphuric  acid.1  By 
substituting  the  two  hydroxyl  groups  of  pyrosulphuric  acid  with  chlorine 
atoms,  pyrosulphuric-acid-chloride  is  obtained  which  is  stable  and  easy  to 
obtain  pure.  (See  No.  150.) 

CARD'S  ACID,  HO  .  SO2  .  OOH,  is  formed  by  the  withdrawal  of  a  molecule 
of  water  from  between  a  molecule  of  sulphuric  acid  and  a  molecule  of  hydro- 
gen peroxide. 

PERSULPHURIC  ACID,  H2S2O8  (see  No.  94),  is  produced  by  the  electrolysis  of 
H2SO4  when  two  discharged  monovalent  acid-sulphate  anions,  HSO4,  become 
united  ;  it  may  also  be  regarded  as  resulting  from  the  condensation  of  one  mole- 
cule of  sulphuric  acid  with  one  of  Caro's  acid. 

THIOSULPHURIC  ACID.  Just  as  sulphuric  acid  may  be  formed  by  the 
addition  of  an  atom  of  oxygen,  so  thiosulphuric  acid,  H2S2OH,  is  produced  by  the 
addition  of  an  atom  of  sulphur  to  a  molecule  of  sulphurous  acid  (cf.  No.  90). 

POLYTHIONIC  ACIDS.  Among  the  polythionic  acids  are  included  dithionic 
acid,  H2S2OG,  trithionic  acid,  H2S3O3,  tetrathionic  acid,  H^Og,  and  pentathionic 


1  Concerning  the  various,  hydrates  of  SO3,  see  R.  Knietsch,  Ber.  34,  4100 
(1901).     Notice  particularly -the  curves.  . 


128 


OXY-ACIDS  OF  SULPHUR. 


acid,  H2S5Ofl.  These  can  be  prepared  only  in  aqueous  solutions  or  in  the 
form  of  their  salts.  (Cf.  Nos.  91,  Barium  Dithionate,  and  92,  Sodium  Tetra- 
thionate.)  The  acids  readily  change  into  one  another,  and  their  constitution 
Is  doubtful. 

HYPOSULPHUROUS  ACID,  H2S2O4.     See  No.  93. 

86.   Sulphuric  Acid  from  Pyrite  by  the  Chamber  Process.1 

Construct  the  apparatus  shown  in  Fig.  20.  The  wide-mouthed 
liter  bottle  is  closed  with  a  cork  through  which  three  holes  are 
bored.  In  one  hole  the  bent-side  arm  of  a  25  c.c.  distilling  flask  is 


FIG.  20. 

inserted.  A  delivery  tube  is  passed  through  the  second  boring  to 
the  middle  of  the  bottle  and  on  the  outside  is  joined  to  a  short 
combustion  tube.  In  the  third  boring  a  tube  is  inserted  which 
leads  through  the  safety  bottle  to  a  suction  pump.  Charge  the 
combustion  tube  with  about  10  g.  of  powdered  pyrite,  and  place  a 
small,  loose  plug  of  asbestos  just  beyond  the  powder  to  retain  any 
unburnt  sulphur.  Place  about  10  c.c.  of  concentrated  nitric  acid 
in  the  distilling  flask,  and  insert  a  glass  tube  through  the  stopper  so 
that  the  lower  end,  which  is  drawn  out  to  a  capillary,  reaches  to 
the  bottom.  Half  fill  the  safety  bottle  with  water  to  show  the 

1  From  Alexander  Smith  and  William  J.  Hale's  A  Laboratory  Outline  of 
General  Chemistry. 


ANTIMONY   SULPHATE.  129 

rate  at  which  air  is  drawn  through  the  apparatus.  The  total  air 
admitted  is  regulated  by  a  screw  clamp  placed  between  the  safety- 
bottle  and  the  water-pump,  and  the  proportion  passing  over  the 
pyrite  and  through  the  nitric  acid  is  governed  by  the  clamp  shown 
at  the  end  of  the  combustion  tube. 

First  heat  the  pyrite  in  a  slow  stream  of  air  until  it  takes  fire. 
Then  warm  the  nitric  acid  and,  by  partially  closing  the  clamp  at 
the  end  of  the  "  pyrite-burner,"  cause  air  to  enter  also  through  the 
capillary  in  the  distilling  flask,  and  thus  carry  nitric  acid  vapors 
into  the  large  bottle.  Heat  the  pyrite  strongly  and  continuously; 
regulate  the  current  of  air  laden  with  nitric  acid  so  that  red  fumes 
are  always  present  in  the  bottle.  If  insufficient  nitric  acid  is  pro- 
vided, the  walls  of  the  flask  become  coated  with  colorless  crystals 
of  nitrosyl  sulphuric  acid  ("  chamber  crystals  ";  cf.  No.  152). 

When  the  pyrite  is  completely  burned,  disconnect  the  apparatus, 
and  wash  the  contents  of  the  bottle  into  a  beaker  with  a  little 
water.  Test  the  solution  qualitatively,  and  determine  by  titration 
the  yield  of  sulphuric  acid. 

87.  Reduction  of  Barium  Sulphate  and  Preparation  of  Barium 

Nitrate. 

Mix  thoroughly  47  g.  of  finely  powdered  heavy  spar  with  12  g. 
of  fine  sifted  charcoal,  and  place  the  mixture  in  a  Hessian  crucible 
which  should  be  two-thirds  to  three-quarters  filled  by  it.  Spread 
a  layer  of  charcoal  on  top,  and  place  a  cover  on  the  crucible. 
Heat  to  a  red  heat  for  two  hours  in  a  charcoal  furnace.  After 
cooling  break  up  the  reddish-gray,  porous  contents  of  the  crucible, 
and  add  it  little  by  little  to  600  c.c.  of  a  solution  containing  25  g. 
of  nitric  acid.  Boil,  filter,  and  evaporate  the  solution  to  crystalli- 
zation. Separate  the  crystals  on  a  filter,  wash  them  with  a  little 
water,  and  dry  them  at  a  moderate  temperature  on  a  porous  plate. 
Recover  more  crystals  from  the  mother-liquor.  Yield,  45  to  50  g. 
of  Ba(NO3)2. 

88.   Antimony  Sulphate  Sb2(SO4)3. 

Antimony  sulphate  is  an  example  of  a  salt  which  is  very  easily  hydrolyzed. 
In  spite  of  the  fact  that  it  can  be  crystallized  from  concentrated  sulphuric 
acid,  it  is  impossible  to  obtain  it  pure  without  adopting  some  special  expe- 
dient. The  sulphuric  acid  adhering  to  the  crystals  cannot  be  removed  by 


130  OXY-ACIDS  AND  SALTS  OF  SULPHUR. 

suction  or  by  evaporation.  Washing  with  anhydrous  acetic  acid  is  effective, 
however,  because  this  solvent  does  not  react  appreciably  with  either  sulphuric 
acid  or  antimony  sulphate,  and  the  excess  can  be  easily  removed. 

Add  10  g.  of  finely  powdered  antimony  (from  No.  7  or  10)  in 
small  portions  to  100  c.c.  of  hot,  concentrated  sulphuric  acid,  and 
heat  the  mixture  in  a  beaker  nearly  to  boiling  until  all  the  anti- 
mony has  dissolved.  On  cooling,  antimony  sulphate  crystallizes 
in  long,  colorless,  lustrous  needles.  Allow  the  solution  to  cool 
completely,  dilute  it  with  80  c.c.  of  anhydrous  acetic  acid,  and 
again  permit  the  mass  to  cool.  Drain  the  crystals  by  suction  on 
a  hardened  filter,  and  wash  them  rapidly  with  a  little  anhydrous 
acetic  acid  and  then  with  ether;  allow  the  product  to  dry  for  one 
or  two  days  in  a  vacuum  desiccator  over  sulphuric  acid.  Yield, 
13  g. 

If  a  sample  of  the  salt  is  treated  with  water,  it  undergoes 
hydrolysis,  and  free  sulphuric  acid  is  found  in  the  solution  decanted 
from  the  insoluble  basic  salt. 

For  the  quantitative  analysis  of  the  salt,  dissolve  a  weighed 
sample  in  a  little  concentrated  hydrochloric  acid,  add  tartaric 
acid,  dilute  well,  and  precipitate  with  barium  chloride.  From  a 
second  solution  prepared  in  the  same  way,  precipitate  the  antimony 
with  hydrogen  sulphide  and  weigh  it  in  a  Gooch  crucible  as  Sb2S3 
after  drying  it  in  an  atmosphere  of  carbon  dioxide  at  280°. 

89.   Alum  from  Kaolin. 

Potassium  aluminium  sulphate,  K2SO4 .  A12(SO4)3. 24  H2O,  is  distin- 
guished from  other  aluminium  salts  by  its  marked  ability  to  crystallize,  and 
on  this  account  it  can  readily  be  obtained  pure.  Thus  it  has  long  been  the 
most  important  salt  of  aluminium  in  spite  of  the  fact  that  it  contains  but 
5.71%  of  that  element.  More  recently,  however,  aluminium  sulphate  has 
been  used  a  great  deal  instead  of  alum;  it  is  prepared  by  neutralizing  with  sul- 
phuric acid  the  hydrated  aluminium  oxide,  which  can  now  be  prepared  pure 
from  the  mineral  bauxite. 

Stir  50  g.  of  powdered  kaolin  with  75  g.  of  concentrated  sul- 
phuric acid  in  an  evaporating  dish;  heat  the  mixture  for  2^  hours 
on  a  Babo  funnel,  at  first  gently  and  then  more  strongly,  until 
white  fumes  escape.  Triturate  the  mass  when  cold,  and  extract 
once  with  300  c.c.  and  then  three  times  with  100  c.c.  of  boiling 
water;  the  last  filtrate  should  show  no  test  for  aluminium  with 


BARIUM  DITHIONATE.  131 

ammonia.  Dissolve  28  g.  of  potassium  sulphate  (15%  less  than 
the  calculated  amount)  in  the  combined  filtrates,  and  bring  the 
double  salt  to  crystallization  in  the  usual  way.  Recrystallize  the 
combined  crops  of  crystals.  Yield,  about  150  g. 

90.  Sodium  Thiosulphate,  Na2S2O3  -  5  H2O. 

Add  16  g.  of  finely  powdered  roll  sulphur  (not  flowers  of  sulphur) 
to  a  solution  of  126  g.  crystallized  sodium  sulphite,  Na2SO3»7  H2O, 
in  250  c.c.  of  water,  and  boil  the  mixture  in  a  flask  until,  at  the 
end  of  about  two  hours,  the  sulphur  has  dissolved.  Filter,  evapo- 
rate to  the  point  of  crystallization,  and  concentrate  the  mother- 
liquor  to  obtain  further  crops  of  crystals.  The  yield  is  nearly 
quantitative. 

The  above  reaction,  Na2S03+ S  <=»  Na2S2O3,  is  N  reversible. 
Heat  2-3  g.  of  sodium  thiosulphate  with  about  5  g.  of  copper 
powder  in  a  test-tube  so  that  the  thiosulphate  melts  in  its  water 
of  crystallization  and  the  latter  partly  boils  away;  the  copper 
becomes  black  (copper  sulphide),  and  the  aqueous  extract  when 
treated  with  oalcium  chloride  gives  a  precipitate  of  calcium 
sulphite;  calcium  thiosulphate  is  easily  soluble. 

91.  Barium  Dithionate,  BaS2O0  •  2  H2O. 

Prepare  sulphur  dioxide  by  allowing  sulphuric  acid  to  drop  into 
160  c.c.  of  commercial  bisulphite  liquor  (see  note  on  p.  71),  and 
pass  the  gas  into  a  suspension  of  50  g.  of  finely  powdered  pyrolu- 
site  in  250  c.c.  of  water  until  the  latter  is  saturated  and  nearly 
all  of  the  black  manganese  dioxide  has  dissolved.  During  this 
reaction,  which  lasts  about  two  hours,  keep  the  mixture  cooled 
with  water. 

Dilute  to  1.5  liters,  bring  the  solution  to  boiling,  and  keep  it  at 
this  temperature  while  adding  a  solution  of  200  g.  of  crystallized 
barium  hydroxide.  Filter,  and  test  the  filtrate  with  ammonium 
sulphide  for  manganese.  If  the  manganese  is  shown  to  be  all  pre- 
cipitated, remove  the  excess  of  barium  hydroxide  from  the  filt- 
rate by  passing  carbon  dioxide  into  the  boiling  solution.  Filter, 
evaporate  to  one-half,  filter  again,  and  allow  the  salt  to  crystallize. 
Drain  the  crystals  and  evaporate  the  mother-liquor  to  obtain 
another  crop.  Yield,  about  120  g. 


132  OXY-ACIDS  AND  SALTS  OF  SULPHUR. 

A  sample  of  the  salt  dissolves  clear  in  water  and  remains 
unchanged  upon  the  addition  of  a  little  nitric  acid.  By  boiling 
with  nitric  acid,  however,  it  is  oxidized  to  sulphate,  and  oxides  of 
nitrogen  are  evolved. 

92.   Sodium  Tetrathionate,  Na2S4O6  •  2  H2O. 

Triturate  50  g.  of  sodium  thiosulphate,  26  g.  of  iodine,  and  5  g. 
of  water  in  a  mortar  to  a  bright  brownish-yellow  paste.  After  a 
short  time  rinse  the  mass  with  50  c.c.  of  alcohol  into  an  Erlenmeyer 
flask.  At  the  end  of  about  three  hours  drain  the  precipitated 
sodium  tetrathionate  and  wash  it  with  alcohol  until  the  washings 
are  free  from  iodine. 

Dissolve  the  crude  product  in  20  to  25  c.c.  of  lukewarm  water, 
and  by  adding  alcohol  in  portions  of  10  c.c. — in  all  50  c.c. — 
bring  about  a  separation  of  crystals.  After  about  ten  hours, 
during  which  time  the  mixture  has  stood  out  of  contact  with  the 
air  in  an  Erlenmeyer  flask,  or  in  a  vacuum  desiccator,  drain  the 
crystals,  wash  with  alcohol,  and  dry  in  a  desiccator  over  sulphuric 
acid.  The  yield  is  about  20  g.  of  compact,  colorless,  crystalline 
aggregates. 

Reactions:  A  solution  prepared  in  the  cold  gives  no  precipitate 
with  copper  sulphate  even  on  boiling;  with  mercuric  nitrate  a 
yellow  precipitate  is  obtained  which  becomes  black  on  boiling; 
with  mercuric  chloride  a  yellow  flocculent  precipitate  separates 
slowly.  Compare  the  behavior  of  sodium  thiosulphate  and 
sodium  sulphite  with  these  reagents. 

Analysis:  Determine  the  water  of  crystallization  in  a  0.7—1.0  g. 
sample  by  heating  it  to  constant  weight  in  the  steam  closet;  deter- 
mine the  sulphur  by  dissolving  in  water,  oxidizing  with  bromine, 
and  precipitating  with  barium  chloride.  The  amount  of  water 
contained  in  the  salt  prepared  in  this  way  is  stated  variously  in 
the  literature. 

93.    Hyposulphurous  Acid,  H2S2O4. 

Zinc  dissolves  in  aqueous  sulphurous  acid  without  evolution  of  hydrogen, 
and  forms  a  yellow  solution  of  great  reducing  power.  By  studying  the 
properties  of  this  solution  and  later  by  preparing  the  solid  sodium  salt,  the 
formula  of  free  hyposulphurous  acid  was  established  as  H2S2O4. 


POTASSIUM  PERSULPHATE. 


133 


The  older  name  for  hyposulphurous  acid  is  hydrosulphurous  add.  It 
should  not  be  confused  with  thiosulphuric  acid,  H2S2O3,  the  sodium  salt  of 
which  is  commonly  known  as  "hypo." 

Take  two  samples  of  the  same  solution  of  sulphurous  acid,  allow 
one  of  them  to  stand  a  short  time  in  contact  with  a  zinc  rod,  and 
afterward  test  the  reducing  power  of  each  solution  towards  dilute 
indigo. 

94.   Potassium  Persulphate,  Electrolytically. 

In  concentrated  solutions  of  acid  sulphates,  dissociation  takes  place  for 
the  most  part  only  partially: 

KHSO4  <=»  K+  +  HSO4~. 

On  electrolyzing,  and  particularly  with  high-current  densities,  the  acid  sul- 
phate ions  on  becoming  discharged  at  the  anode  unite  in  pairs  to  form  per- 
sulphuric  acid,  H2S2O8,  the  potassium  salt  of  which  is  insoluble. 

Inside  a  large  beaker  filled  with  ice  water,  place  a  glass  cylinder, 
or  a  small  beaker  14  cm.  high  and  6  cm.  in  diameter.  Suspend  in 
this,  by  means  of  a  wire  triangle,  an  11  cm.  long,  2.7  cm.  wide 
glass  tube  open  at  both  ends  (a  test-tube  with  its  bottom  cut  off). 
Use  for  the  cathode  a  loop  of  platinum  wire, 
placed  as  near  the  surface  of  the  solution  as 
possible  and  outside  the  inner  tube;  for  the 
anode  melt  a  platinum  wire  into  a  glass  tube 
so  that  it  projects  1.5  to  2.0  cm.,  and  insert 
this  through  the  inner  tube  until  it  reaches 
nearly  to  the  bottom  of  the  beaker  (see  Fig.  21). 

Fill  the  inside  beaker  one-half  full  of  a  satu- 
rated solution  of  acid  potassium  sulphate.  Use 
a  current  density  at  the  anode  of  100  amperes 
per  100  sq.  cm.  of  electrode  surface.  Measure 
the  protruding  anode  wire  and  estimate  its 
surface.  The  current  will  amount  to  less  than 
an  ampere  with  a  moderately  stout  platinum 
wire.  The  temperature  of  the  electrolyte  should  be  lower  than 
15°  and  may  quite  easily  be  maintained  at  6°  to  8°.  A  few  min- 
utes after  closing  the  circuit,  crystals  of  the  difficultly  soluble 
persulphate  are  seen  to  separate.  Continue  the  electrolysis  40 
minutes  in  one  experiment  and  an  hour  in  another.  Collect  the 
salt  on  a  hardened  filter,  wash  it  with  alcohol  and  then  with  ether. 
Weigh  the  salt,  after  drying  in  a  desiccator,  and  analyze  it  as  soon 


FIG.  21. 


134  CARBONATES. 

as  possible  in  the  following  manner:  Dissolve  0.3  g.  of  the  salt  in  a 
beaker  with  10  c.c.  of  0.1-normal  ferrous  ammonium  sulphate  l 
solution  and  200  c.c.  of  hot  water.  After  cooling,  titrate  back  the 
unoxidized  ferrous  salt  with  0.1-normal  permanganate. 

According  to  the  above  directions,  1.25  g.  of  the  persulphate  of 
92  to  95%  purity  should  be  obtained  in  40  minutes  and  1.65  g.  in 
an  hour,  and  the  current  yield  should  be  from  41  to  43%. 

To  test  qualitatively  the  oxidizing  power  of  the  salt,  dissolve 
some  of  it  in  water,  make  the  solution  alkaline,  add  it  to  a  solution 
of  a  lead  or  manganese  salt  to  which  alkali  has  also  been  added, 
and  warm  the  mixture. 

(/)    Carbonates. 
95.    Sodium  Carbonate  (Ammonia-Soda  Process). 

From  a  concentrated  solution  containing  the  ions  Na+,  Cl~,  NH4+,  and 
HCO3~~,  sodium  bicarbonate,  NaHCO3,  the  most  insoluble  of  the  possible 
salts,  is  the  first  to  separate  (cf.  No.  80).  The  acid  salt  loses  carbon  diox- 
ide when  heated  and  changes  into  the  monocarbonate,  Na2CO3.  This  pro- 
cess, which  was  first  placed  on  a  commercial  basis  by  Solvay,  yields  at  the 
present  time  the  largest  part  of  the  world's  soda.  For  a  technical  discussion 
of  the  process,  consult  Lunge,  Sulphuric  Acid  and  Alkali,  Vol.  Ill;  for  a 
theoretical  treatment  in  the  light  of  the  phase  rule,  see  Bodlander  and  Breull, 
Z.  angew.  Chem.  14,  381,  405  (1901);  and  P.P.  Fedotieff,  Z.  physikal.  Chem. 
49,  162  (1904). 

Add  60  g.  of  pulverized  sodium  chloride  to  180  g.  of  a  10% 
ammonia  solution,  and  allow  it  to  stand  with  occasional  shaking 
until  nearly  all  has  dissolved.  Place  the  filtered  solution  in  a  closed 
flask  and  saturate  it  at  room  temperature  with  carbon  dioxide;  pass 
the  gas  through  a  wash  bottle  containing  water  and  then  allow  it 
to  enter  the  flask  through  a  single  inlet  tube,  which  dips  under 
the  solution,  as  rapidly  as  it  will  be  absorbed  (in  all,  perhaps  24 
hours).  Collect  the  precipitated  sodium  bicarbonate  on  a  filter 
and  wash  it  with  a  little  cold  water.  Dry  the  product  in  a 
porcelain  dish  and  then  heat  it  with  a  free  flame  until  carbon 
dioxide  ceases  to  escape.  Recrystallize  the  crude  sodium  carbonate 
thus  obtained  from  five  times  its  weight  of  water,  and  wash  the 
crystals  with  a  little  water.  Concentrate  the  mother-liquor  to 
obtain  another  crop  of  crystals.  Recrystallize.  Yield,  about  40  g. 

1  40  g.  of  (NH4)2SO4.FeSO4.6H2O  and  30  g.  of  cone.  H2SO4  made  up  to 
1  liter  and  standardized  against  0.1-normal  permanganate. 


BARIUM  HYPOPHOSPHITE.  135 

A  solution  of  the  sodium  carbonate,  after  being  acidified  with 
nitric  acid,  ought  not  to  show  more  than  a  slight  cloudiness  with 
silver  nitrate. 

(g)    Phosphoric  Acids.1 

Structural  formulas  can  be  written  as  below  for  orthophosphoric  acid, 
H3PO4,  phosphorous  acid,  H3PO3,  and  hypophosphorous  acid,  H.<PO2,  if  it  is 
taken  into  consideration  that  the  first  acid  is  tribasic,  the  second  dibasic  and 
the  third  monobasic,  and  if  it  is  assumed,  although  this  is  by  no  means  proved, 
that  all  the  hydroxyl  hydrogens,  and  only  these,  can  be  replaced  by  metals 
in  the  formation  of  salts. 


Orthophcsphoric  acid.        Phosphorous  acid.  Hypophosphorous  acid. 

The  reducing  action  of  the  last  two  acids  would  accordingly  consist  in 
taking  up  oxygen  which  would  enter  between  the  phosphorus  atom  and  the 
hydrogen  atoms. 

If  it  is  not  assumed  that  all  the  hydroxyl  hydrogens  are  replaceable,  then 
the  formulas  P(OH)5  and  HP(OH)2  become  possit^je,  according  to  which  phos- 
phorous and  hypophosphorous  acids  are  derived  from  trivalent  phosphorus. 
It  may  be,  as  in  the  case  of  sulphurous  acid,  that  both  formulas  have  a  cer- 
tain justification  (compare  p.  126  and  Nos.  155  and  156). 

By  the  loss  of  water  from  one  and  from  two  molecules  of  orthophosphoric 
acid,  meta-phosphoric  and  pyro-phosphonc  acids  respectively  are  obtained: 

(  -  O  O  =  )  (  =  O 

P  }  =  O  HO->P-O-P}-OH 

{  -  OH  HO  -  $  (  -  OH 

Metaphosphoric  acid.  Pyrophosphoric  acid. 

HYPOPHOSPHORIC  ACID,  H2PO3,  which  was  formerly  written  H4P2O6,  has 
recently  been  found  to  have  a  molecular  weight  corresponding  to  the  smaller 
molecule,  and  is  presumably  derived  from  tetravalent  phosphorus. 

96.   Barium  Hypophosphite,  Ba(H2PO2)2 .  H2O. 

Salts  of  hypophosphorous  acid  are  formed  by  the  action  of  phosphorus 
on  warm  aqueous  solutions  of  strong  bases: 

4  P  +  3  KOH  +  3  H2O  =  3  KH2PO2  +  PH3. 

If  barium  hydroxide  is  chosen  as  the  base,  it  is  possible  to  obtain  the  crys- 
tallized salt,  after  precipitating  the  excess  of  the  base  with  carbon  dioxide 
and  evaporating  the  filtrate.  Hypophosphites  are  strong  reducing  agents. 

1  Cf.  Note  2,  p.  126. 


136  THIO-ACIDS  AND  SALTS. 

Heat  a  solution  of  120  g.  crystallized  barium  hydroxide  in 
1200  c.c.  of  water  with  30  g.  of  yellow  phosphorus  (Caution)  in 
a  round-bottomed  flask  on  a  Babo  funnel.  After  about  four  hours 
nearly  all  of  the  phosphorus  will  have  disappeared.  The  process 
should  be  carried  out  in  a  well-ventilated  hood,  since  phosphine 
gas  escapes  freely.  Filter  the  solution  through  a  plaited  filter 
into  a  large  porcelain  dish;  heat  and  pass  in  carbon  dioxide  until 
the  excess  of  barium  hydroxide  is  precipitated.  Filter,  rinse  the 
precipitate  with  boiling  water,  and  evaporate  the  solution  to  one- 
half  its  volume.  After  again  filtering,  concentrate  further  —  at 
the  last  in  a  beaker  —  until  crystals  begin  to  separate ;  then  add 
alcohol  and  leave  the  solution  to  crystallize.  Collect  the  product 
on  a  filter  and  evaporate  the  mother-liquor  to  obtain  more  crystals. 
Purify  the  entire  product  by  recrystallization.  Yield,  40-60  g.  of 
colorless  flaky  crystals. 

Reactions:  1.  Treat  a  sample  of  the  product  with  concentrated 
sulphuric  acid  and  heat  it  to  boiling;  a  large  amount  of  sulphur 
dioxide  escapes  and  sulphur  distils  to  the  cooler  part  of  the  test-tube. 

2.  On  adding  hypophosphite  to  a  dilute  solution  of  gold  chloride, 
H[AuCl4],  and  warming  gently,  a  blue-violet  coloration  (cf.  No.  25) 
and  later  a  violet-red  precipitate  of  gold  is  obtained. 

3.  With  silver  nitrate  a  dark  brown  separation  of  silver  takes 
place,  slowly  at  the  room  temperature  and  more  rapidly  when 
warm;  with  mercurous  salts  a  separation  of  mercury  occurs. 

4.  If  a  dry  sample  of  barium  hypophosphite  is  heated  in  a  test- 
tube,  it  turns  red  with  loss  of  water  and  further  decomposition: 
phosphorus  distils  off  and  phosphine  escapes. 

Dependent  preparation:  Copper  Hydride,  No.  33. 

THIO-ACIDS  AND  THEIR  SALTS. 

97.   Potassium  Trithiocarbonate  Solution  (Reagent  for  Nickel). 

Divide  50  c.c.  of  a  5%  solution  of  potassium  hydroxide  in  two 
equal  portions;  saturate  one  with  hydrogen  sulphide  and  mix  it 
with  the  other  portion.  Shake  this  solution  of  potassium  sulphide 
vigorously  with  2  c.c.  of  carbon  bisulphide  for  five  minutes  and 
then  pour  it  through  a  filter  which  has  been  moistened  with  water. 
The  bright,  orange-red  solution  contains  K2CS3,  and  is  an  extremely 
delicate  reagent  for  nickel.  With  concentrated  ammoniacal  solu- 


SODIUM  THIOANTIMONATE.  137 

tions  of  nickel  salts,  it  gives  a  brownish-black  precipitate;  with 
dilute  solutions  a  dark  brown  coloration. 

Determine  the  sensitiveness  of  the  test:  Start  with  1  or  2  c.c.  of 
a  0.0002-normal  nickel  solution,  and  if  this  gives  a  distinct  reaction, 
dilute  ten  times  and  test  again,  and  continue  in  this  way  until  the 
limit  of  sensitiveness  is  reached. 

In  order  to  find  to  what  extent  cobalt  interferes  with  the  reaction, 
test  solutions  containing  1,  2,  10,  50  and  more  atomic  equivalents 
of  cobalt  for  each  atom  of  nickel. 

98.    Barium  Trithiocarbonate. 

Dissolve  32  g.  of  crystallized  barium  hydroxide  in  100  c.c.  of  hot 
water.  Place  one-half  of  the  solution  in  a  closed  flask  and  saturate 
it  with  hydrogen  sulphide,  whereby  barium  sulphydrate  is  formed; 
then  add  the  other  half  of  the  original  barium  hydroxide  solution. 
Shake  the  resulting  solution  of  barium  sulphide  with  8  g.  of  carbon 
bisulphide,  whereupon  barium  trithiocarbonate,  BaCS3,  precipi- 
tates as  a  yellow,  crystalline  powder.  Drain  the  precipitate,  wash 
it  with  a  little  water,  then  with  50%  alcohol,  and  finally  with  pure 
alcohol;  and  dry  it  in  a  warm  place  (on  top  of  the  hot  closet).  As 
the  alcohol  used  for  washing  runs  into  the  filtrate,  more  of  the 
barium  trithiocarbonate  precipitates.  Yield,  12-15  g. 

99.    Sodium  Thioantimonate,  Na3SbS4 . 9  H2O. 

Sodium  thioantimonate,  or  "Schlippe's  salt,"  is  formed  by  the  inter- 
action of  antimony  sulphide,  sodium  sulphide,  and  sulphur;  sodium  thio- 
antimonite  may  first  be  formed  in  the  dry  way  (Schlippe),  or  the  whole 
process  may  be  carried  out  in  the  wet  way  (Mitscherlich).  According  to 
Schlippe's  method,  sodium  sulphate  is  heated  together  with  charcoal  and 
antimony  trisulphide  in  the  furnace;  sodium  sulphide  is  formed  by  reduc- 
tion, and  this  immediately  combines  with  the  antimony  sulphide.  By 
suspending  the  pulverized  fusion  in  water -and  boiling  it  with  powdered 
sulphur,  the  sodium  thioantimonite  is  converted  into  thioantimonate. 

In  recrystallizing  the  salt,  it  is  necessary  to  add  a  little  sodium  hydroxide 
to  prevent  hydrolysis,  for  otherwise  the  free  thioantimonic  acid  which  would 
form  would  decompose  into  antimony  sulphide  and  hydrogen  sulphide. 
Sodium  thioantimonate  is  used  in  making  the  medicinal  preparation  of 
antimony  pentasulphide. 

1.  Mix  thoroughly  36  g.  of  powdered  stibnite  (34  g.  =  0.1  mol.), 
43  g.  of  anhydrous  sodium  sulphate,  and  16  g.  of  powdered  charcoal. 
Place  the  mixture  in  a  Hessian  crucible,  which  should  be  about  half 


138  THIOACIDS   AND  SALTS. 

filled,  cover  with  a  layer  of  charcoal,  and  heat  in  a  charcoal  furnace 
until  the  charge  comes  to  a  state  of  quiet  fusion,  and  then  heat  ten 
minutes  longer.  Pour  the  melt  upon  an  iron  plate,  pulverize  the 
mass  when  cold,  and  boil  it  half  an  hour  with  7  g.  of  flowers  of  sul- 
phur and  300  c.c.  of  water.  Add  a  little  caustic  soda  solution  and 
evaporate  the  nitrate  to  crystallization  in  a  porcelain  dish.  Collect 
the  crystals,  wash  them  with  a  little  alcohol,  and  work  up  the 
mother-liquor  further.  Recrystallize  the  combined  portions,  add- 
ing a  little  caustic  soda  to  the  solution;  dry  the  product  in  a  vacuum 
desiccator  over  lime,  upon  which  a  few  drops  of  ammonium  sul- 
phide have  been  poured.  The  yield  is  40  to  50  g.  of  light-yellow 
well-formed  crystals. 

2.  Slake  26  g.  of  quicklime  with  hot  water,  stir  it  up  to  a  paste 
with  an  additional  80  c.c.  of  water,  and  add  a  solution  of  70  g. 
crystallized  sodium  carbonate  in  250  c.c.  of  water.  Bring  the 
mixture  to  boiling  in  an  iron  dish,  and  while  it  is  boiling  add  little 
by  little  a  paste  made  from  36  g.  of  powdered  stibnite,  7  g.  of  pow- 
dered sulphur,  and  12  g.  of  water.  When,  after  boiling  about 
15  minutes,  the  gray  color  of  the  antimony  sulphide  has  disap- 
peared, filter  the  solution  through  linen  cloth  and  extract  the 
residue  by  boiling  it  with  100  to  150  c.c.  of  water.  Bring  the 
combined  filtrates  to  crystallization  and  proceed  as  in  Method  1 
with  the  recrystallization,  etc.  Yield,  about  60  g.1 

100.   Potassium  Ferric  Sulphide,  K[FeS2]. 

Place  an  intimate  mixture  of  about  30  g.  iron  powder,  180  g. 
flowers  of  sulphur,  150  g.  potassium  carbonate,  and  30  g.  anhydrous 
sodium  carbonate  in  a  Hessian  crucible,  and  heat  the  mass  in  a 
charcoal  furnace  until  it  is  melted  to  a  thin  liquid.  This  takes 
about  an  hour.  Close  the  furnace  and  let  the  melt  cool  slowly; 
break  the  crucible  and  digest  the  lumps  of  the  melt  with  warm 
water  in  a  porcelain  dish  until  they  are  completely  disintegrated. 
From  time  to  time  replace  the  resulting  green  solution  with  fresh 
water  until  nothing  more  dissolves  and  pure,  glistening,  dark 
needles  remain  behind.  Wash  the  product  with  water  and  alcohol 
and  dry  it  in  the  steam  closet.  Yield,  about  70  g.  Confirm  the 
composition  by  a  quantitative  analysis. 

1  The  yield  corresponds  closely  with  the  reaction  as  given  by  Mitscherlich : 
8  Sb2S3  +  16  S  +  18  Na2CO3  +  3  H2O  =  10  Na3SbS4  +  3  H2Na2Sb2Or  +  18  CO2. 


COMPLEX  HALOGEN  ACIDS  AND  SALTS.  139 

101.   Ammonium  Copper  Tetrasulphide,  NH4[CuSj. 

It  is  known  from  qualitative  analysis  that  copper  sulphide  dissolves  appre- 
ciably in  ammonium  or  sodium  polysulphide.  Alkali  salts  of  thio-copper  acids 
are  thereby  formed  of  which  the  one  under  consideration  has  been  best  in- 
vestigated.1 

Place  a  mixture  of  200  c.c.  of  concentrated  ammonia  and  50  c.c. 
of  water  in  a  closed  flask,  and  while  keeping  the  liquid  cooled  with 
tap  water,  pass  in  hydrogen  sulphide  until  it  is  saturated.  Dis- 
solve in  one-half  of  the  solution  as  much  finely  powdered  sulphur 
as  possible  at  40°  (about  60  g.),  then  filter  this  and  add  it  to  the 
other  half  of  the  solution. 

While  rotating  this  solution  in  a  flask,  add  a  10%  solution  of  blue 
vitriol  little  by  little  until  a  permanent  precipitate  of  copper  sul- 
phide just  begins  to  form;  filter  immediately  through  a  plaited 
filter  into  an  Erlenmeyer  flask;  the  latter  should  be  filled  almost 
completely  with  the  liquid.  On  standing,  best  in  an  ice-box, 
brilliant  red  prisms  separate,  which  on  the  next  day  should  be 
washed  with  water  and  then  with  alcohol  and  dried  quickly  over 
lime  in  a  vacuum  desiccator.  On  adding  copper  sulphate  to  the 
main  filtrate  (to  which  the  washings  should  not  have  been  added) 
a  further  yield  of  the  crystals  is  obtained.  The  entire  product  is 
about  25  g. 

A  sample  of  the  salt  dissolves  completely  in  a  small  amount  of 
2-normal  sodium  hydroxide,,  which  points  to  the  existence  of  the 
copper  as  a  part  of  the  complex;  after  some  time  copper  sulphide 
begins  to  separate  slowly.  Ammonia  is  readily  detected  in  this 
solution  by  its  odor  or  by  Nessler's  reagent.  Concentrated  potas- 
sium hydroxide  added  to  the  fresh  solution  immediately  pre- 
cipitates red  potassium  copper  tetrasulphide.  Concentrated 
hydrochloric  and  nitric  acids  act  but  slowly  on  the  dry  salt,  con- 
centrated sulphuric  acid  not  at  all. 

COMPLEX  HALOGEN  ACIDS  AND  THEIR  SALTS.  — COM- 
PLEX  CYANOGEN  COMPOUNDS. 

Among  the  complex  compounds  formed  by  the  union  of  two  simple  sub- 
stances, the  particular  class  in  which  both  of  the  simple  substances  are 
halogen  or  cyanogen  compounds  requires  a  special  treatment,  —  first,  for 
the  reason  that  the  number  of  such  compounds  is  very  great,  and  many  of 
them  are  of  considerable  practical  importance;  and  second,  because  their 

1  H.  Blitz  and  P.  Herms,  Ber.  40,  977  (1907). 


140  COMPLEX  HALOGEN  ACIDS  AND  SALTS. 

classification  and  interpretation  on  the  basis  of  the  old  theories  of  valence 
have  for  a  long  time  caused  chemists  much  perplexity. 

For  the  diagnosis  of  such  complex  compounds,  purely  chemical  tests  are 
first  made  as  to  the  reactivity  of  their  constituents.  If  the  typical  reactions 
of  the  simple  salts  do  not  occur  with  these  double  compounds,  but  instead 
new  specific  reactions  are  found,  then  a  complex  structure  of  the  substance 
is  indicated.  Such  cases  are  abundant  in  analytical  chemistry,  as,  for 
example,  with  the  iron-cyanogen  compounds. 

Further,  the  existence  in  aqueous  solution  of  complex  metal-containing 
ions  can  be  proved  electrochemically,  and  with  this  proof  began  the  clearing 
up  of  the  constitution  of  these  substances.  The  metal  migrates,  as  a  constit- 
uent of  the  anion,  toward  the  positive  pole  and  there  increases  in  concentra- 
tion. Thus  an  anomalous  change  in  concentration  takes  place  with  respect 
to  the  metal  when  it  is  compared  with  that  occurring  in  the  electrolysis 
of  simple  metal  salts.  From  this  Hittorf,  in  the  years  1853-1859,  deduced 
the  existence  of  complex  ions  in  solution.  Previously  Porett  (1814)  had 
noticed,  incidentally,  the  analogy  existing  between  potassium  ferrocyanide 
and  the  salts  of  the  oxy-acids. 

Finally,  the  degree  of  the  complexity  of  dissolved  double  compounds  can 
be  estimated  by  means  of  any  one  of  the  methods  of  determining  ion  concen- 
trations, such  as  by  measurements  of  the  freezing-point,  of  electromotive  force, 
etc.  Thus  not  only  can  the  extreme  cases  be  distinguished  —  in  which  there 
is  merely  a  dissociation  into  a  simple  cation  and  a  complex  anion  or  else  a 
complete  breaking  down  into  the  simple  ions  —  but  also  it  can  be  shown,  in 
the  transition  cases,  to  what  degree  the  two  kinds  of  dissociation  prevail. 

The  extent  to  which  the  nature  of  the  compound,  and  indeed  the  nature 
of  the  elements  forming  the  complex  as  well  as  of  those  outside  of  the  com- 
plex, favors  the  first  or  second  form  of  dissociation  is  now  understood  in  a 
general  way.  The  formation  of  complexes  seems  to  be  favored  when  the 
anions  of  the  simple  compounds  are  the  same.  The  stability  of  the  com- 
plex increases  from  the  chlorine  compounds  through  the  other  halogen  to 
the  cyanogen  compounds;  it  increases  further,  as  far  as  metal-containing 
anions  are  concerned,  with  the  " nobleness"  of  the  metal.  The  latter  property 
of  metals  is  intimately  connected  with  the  discharge  potential  of  their  ions  — 
a  quantity  capable  of  exact  measurement — (electro  affinity,  cf.  p.  59),  hence 
it  would  seem  possible  to  predict,  from  data  concerning  discharge  potentials, 
the  relative  tendency  of  the  metals  to  form  complexes.  Abegg  and  Bodlander 
have  in  fact  proposed  a  new  system  of  chemical  classification  l  based  upon  this 
fundamental  property  of  the  atoms.  Among  the  numerous  investigations 
which  these  ideas  have  instigated,  mention  should  be  made  of  the  measure- 
ments carried  out  by  Bodlander  of  the  so-called  "stability-constants"  which 
give  the  proportion  in  which  simple  metal  ions  dissociate  from  complex  metal- 
containing  anions.2 

1  Abegg  and  Bodlander:   Die   Electroaffinitat.     Z.  anorg.  Chem.  20,  453 
(1899). 

2  Bodlander:  Ber.  36,  3933  (1903). 


HYDROFLUOSILICIC  ACID.  141 

The  composition  of  the  complex  halogen  compounds  shows  that  one 
atom  of  non-haloid  nature  occupies  a  peculiar  position  in  the  complex,  while 
the  remaining  atoms  of  the  complex,  which  are  either  all  of  the  same  kind 
or  closely  related,  are  most  frequently  either  four  or  six  in  number.  Start- 
ing from  these  facts  Werner  has  used  with  success  the  so-called  "Coordina- 
tion Theory,"  *  which  he  himself  proposed,  to  explain  the  constitution  of 
these  compounds.  According  to  this  theory  the  halogen  atoms  are  situated 
in  space  around  the  other  atom  of  the  complex,  the  "central  atom";  when 
they  number  four  they  may  be  supposed  to  be  located  at  the  corners  of  a 
square,  and  when  they  number  six,  at  the  corners  of  an  octahedron  which 
surrounds  the  central  atom  (see  p.  166).  Werner  has,  in  constructing  his 
theory,  also  introduced  the  idea  of  "secondary  valences"  which  differ  frcm 
the  "principal  valences"  by  being  weaker  and  incapable  of  binding  electrons. 
Thus  aluminium,  when  its  three  principal  valences  are  saturated  with  fluo- 
rine, can  bring  into  play  its  three  secondary  valences,  and  thereby  bind  the 
fluorine  atoms  in  three  molecules  of  sodium  fluoride: 


P\         ,-F-]-N 
F-^Afl-F—  —  N 

IF/      XF_  _ N 


1— Na 

Na 
Na 


The  main  advantages  of  Werner's  theory  will  be  shown  more  clearly  in 
specific  cases,  particularly  in  the  chemistry  of  the  complex  cations.  The 
attempts  which  have  been  made  to  represent  structural  formulas  of  these 
substances,  in  accordance  with  the  older  theory  of  valence,  now  possess 
merely  an  historic  interest;  thus  the  formula  of  potassium  ferrocyanide 
has  been  written  as 


J — Fe — C 
or 


K 

but  these  formulas  correspond  neither  to  the  way  the  atoms  are  combined 
in  the  compound,  to  the  dissociation  relations,  nor  to  the  tendency  shown 
by  the  iron,  cyanogen  or  potassium  to  react.  Similar  statements  hold  true 
regarding  many  of  the  structural  formulas  of  the  silicates. 

102.   Hydrofluosilicic  Acid. 

The  reaction,  SiO2  +  4  HF  =  SiF4  +  2  H2O,  is  reversible.  When  warm 
and  in  the  presence  of  sulphuric  acid,  which  has  a  dehydrating  action,  it 
proceeds  from  left  to  right.  If,  however,  silicon  tetrafluoride  is  brought  in 
contact  with  a  large  amount  of  water,  then  silicic  acid  and  hydrofluoric  acid 

1  For  a  full  discussion  of  this  theory  see  A.  Werner,  Neuere  Anschau- 
ungen  auf  dem  Gebiete  der  anorganischen  Chemie,  Braunschweig,  1909. 


142  COMPLEX  HALOGEN  ACIDS  AND  SALTS. 

are  formed,  but  the  hydrofluoric  acid  combines  in  a  secondary  reaction  with 
undecomposed  silicon  fluoride  to  form  hydrofluosilicic  acid: 
2  HF  +  SiF4  =  HJSiFJ. 

Place  a  mixture  of  100  g.  powdered  fluorspar  and  40  g.  of  pre- 
cipitated silicic  acid,  or  80  g.  of  sand,  in  a  round-bottomed  flask, 
and,  while  shaking,  add  500  g.  of  concentrated  sulphuric  acid  in 
small  portions.  Heat  the  mixture  on  a  Babo  funnel  and  conduct 
the  escaping  silicon  fluoride  into  1  liter  of  water  by  means  of  a  dry 
delivery  tube  that  dips  directly  into  mercury  which  one-third  fills 
a  small  beaker  standing  in  the  bottom  of  a  large  beaker  containing 
the  water.  The  lower  end  of  the  delivery  tube  should  be  made 
wider  by  sealing  on  a  short  piece  of  tube  1.5  cm.  in  diameter. 
The  mercury  keeps  the' opening  of  this  tube  dry  and  thus  pre- 
vents it  from  becoming  clogged  with  silicic  acid.  If  the  water 
becomes  too  much  thickened  with  silicic  acid,  remove  a  part  of 
the  liquid  without  interrupting  the  process,  filter  it  through  a 
piece  of  linen  laid  in  a  large  Biichner  funnel,  and  return  the  fil- 
trate to  the  beaker. 

When  the  reaction  is  complete  remove  the  silicic  acid,  as  above, 
with  a  linen  filter,  and  estimate  the  yield  by  analyzing  a  sample  of 
the  filtrate.  Either  titrate  hot  with  0.1-normal  sodium  hydroxide, 
using  phenolphthalein  as  indicator,  whereby  the  reaction  is 

H2SiF6  +  6  NaOH  =  6  NaF  +  H4SiO4  +  2  H20, 
or  else  add  to  the  solution  an  excess  of  neutral  calcium  chloride 
and  titrate  the  hydrochloric  acid  set  free,  using  methyl-orange  as 
indicator: 

H2SiF6  +  3  CaCl2  +  4  H20  =  3  CaF2  +  6  HC1  +  H4SiO4. 

Add  some  potassium  chloride  solution  to  another  sample  of  the 
product;  the  difficultly  soluble  potassium  salt  separates,  and, 
although  the  precipitate  is  barely  visible  at  first,  the  liquid  is 
eventually  left  in  a  jelly-like  condition. 

103.  Potassium  Titanium  Fluoride,  K2[TiFj. 
Allow  19  g.  of  titanium  tetrachloride  (No.  52)  to  flow  very 
slowly  from  a  dropping  funnel  into  20  c.c.  of  ice-cold  water  in  a 
platinum  dish.  (Hood.)  The  dish  must  meanwhile  be  kept  sur- 
rounded with  ice.  Add  30  g.  of  pure,  40-50%  hydrofluoric  acid 
to  the  mixture,  and  then  a  warm  concentrated  solution  of  15  g. 
potassium  chloride;  the  contents  of  the  dish  thereupon  harden 


AMMONIUM  PLUMBIC  CHLORIDE.  143 

to  a  crystalline  paste.  Evaporate  to  complete  dryness  on  the 
water  bath,  to  remove  the  excess  of  hydrofluoric  acid,  and  recrystal- 
lize  the  difficultly  soluble  residue  from  water.  Yield,  about  20  g. 

104.   Ammonium  Plumbic  Chloride,  (NH4),[PbCl6].     Lead 
Tetrachloride,  PbCl4. 

Grind  10  g.  of  lead  chloride,  PbCl2,  with  20  c.c.  of  concentrated 
hydrochloric  acid;  after  allowing  it  to  settle  for  a  short  time, 
decant  the  solution  together  with  the  fine  suspended  solid;  repeat 
the  above  treatment  with  the  residue  until  all  of  the  lead  chloride 
is  in  solution  or  in  a  state  of  finest  suspension  in  200  c.c.  of 
hydrochloric  acid. 

Place  mixtures  prepared  in  this  manner  in  each  of  two  250— 
300  c.c.  gas- wash-bottles,  and  while  keeping  the  temperature  at 
10°-15°  and  shaking  occasionally,  pass  a  slow  current  of  chlorine 
through  the  two  bottles  placed  in  series.  After  about  five  hours 
all  the  lead  chloride  should  be  dissolved.  If  any  remains,  allow 
the  solutions  to  stand  over  night  and  then  filter  off  any  insoluble 
residue  on  asbestos.  Combine  the  two  filtrates  in  a  600  c.c. 
flask,  cool  the  liquid  with  ice,  and  add  an  ice-cold  solution  of  8  g. 
ammonium  chloride  in  80  c.c.  of  water;  after  a  short  time  a  heavy, 
yellow,  crystalline  precipitate  of  ammonium  plumbic  chloride 
begins  to  separate.  After  several  hours  filter  rapidly  on  a  hardened 
filter,  wash  the  precipitate  with  50  c.c.  of  ice-cold  alcohol  and  dry 
it  at  50°.  Yield,  about  20  g. 

On  treating  a  sample  of  the  ammonium  plumbic  chloride  with 
water,  it  decomposes  immediately,  forming  brownish-black  hydrated 
lead  dioxide  which  remains  to  some  extent  in  colloidal  solution. 

Lead  Tetrachloride.  Mix  20  g.  of  ammonium  plumbic  chloride 
thoroughly,  by  means  of  a  mechanically  driven  stirrer,  with  60  c.c. 
of  ice-cold,  concentrated  sulphuric  acid.  In  a  short  time  heavy, 
nearly  colorless,  oily  drops  of  lead  tetrachloride  separate.  By 
repeated  decantation  and  stirring  up  with  fresh  portions  of  sul- 
phuric acid,  it  is  possible  to  separate  all  of  the  ammonium  sulphate 
from  the  oil,  and  the  latter  then  settles  as  a  fairly  clear  layer.  If 
the  oil  is  allowed  to  remain  long  in  contact  with  the  sulphuric 
acid,  lead  sulphate  is  formed. 

Lead  tetrachloride  decomposes  with  explosive  violence  when 
strongly  heated. 


144  COMPLEX  HALOGEN   ACIDS  AND  SALTS 

105.   Potassium  Lead  Iodide,  K[PbI3] .  2  H2O. 

Potassium  lead  iodide  is  stable  only  when  in  contact  with  a  cold,  concen- 
trated solution  of  potassium  iodide.  If  the  solution  is  diluted  with  water, 
or  even  if  it  is  merely  heated,  potassium  iodide  dissolves  out  of  the  solid 
compound,  and  lead  iodide  remains  behind.  On  concentrating  and  cooling 
the  solution,  the  double  salt  is  again  formed. 

If  a  hot  solution  of  4  g.  lead  nitrate  in  15  c.c.  of  water  is  mixed 
with  a  hot  solution  of  15  g.  of  potassium  iodide  in  15  c.c.  of  water, 
yellow  lead  iodide  is  at  first  precipitated.  On  cooling  to  room 
temperature,  the  crystals  of  lead  iodide  disappear  and  a  very  pale- 
yellow,  felted  mass  of  crystal  needles  is  produced.  On  heating, 
the  crystals  of  the  double  compound  disappear  with  a  re-formation 
of  lead  iodide;  on  cooling,  the  double  salt  is  again  produced.  The 
change  can  be  observed  especially  well  if  a  drop  of  the  hot  solution  is 
placed  between  a  heated  slide  and  cover  glass  under  the  microscope. 

If  it  is  desired  to  obtain  the  dry  salt,  it  should  be  collected  on  a 
filter,  and,  without  washing,  pressed  between  filter  papers  and 
dried  in  a  vacuum  desiccator. 

Potassium  lead  iodide  possesses  the  remarkable  property  of 
being  extremely  soluble  in  acetone.  If  the  preparation  is  treated 
while  still  moist  with  10  to  15  c.c.  of  acetone,  there  is  produced, 
even  in  the  cold,  a  yellow  solution  from  which  the  salt  can  be 
precipitated  by  addition  of  two  or  three  volumes  of  ether.  On 
evaporating  the  acetone  solution,  no  well-formed  crystals  are 
obtained.  If  a  few  drops  of  this  solution  are  allowed  to  evaporate 
on  filter  paper  the  salt  is  obtained  in  a  state  of  very  fine  subdivision, 
in  which  condition  it  is  extremely  sensitive  to  the  least  traces  of 
moisture.  Even  the  moisture  of  the  air  suffices  to  decompose 
the  salt  in  a  short  time;  the  yellow  color  of  lead  iodide  which 
thereby  appears  indicates  in  the  sharpest  manner  the  presence  of 
traces  of  water.  If  another  paper  prepared  in  the  same  manner 
is  left  in  a  vacuum  desiccator  over  sulphuric  acid,  no  noticeable 
yellow  color  appears;  on  opening  the  desiccator,  however,  it 
develops  immediately.  * 

106.   Potassium  Mercuric  Iodide,  K2[HgI4] .  2  H2O. 

Precipitate  mercuric  iodide  from  an  aqueous  solution  of  13.5  g. 
mercuric  chloride  by  adding  a  solution  of  16.6  g.  potassium  iodide. 


1 
POTASSIUM   COBALTICYANIDE.  145 

Wash  the  precipitate  and  then  redissolve  it  in  a  hot  solution  of 
16  g.  potassium  iodide  in  10  c.c.  of  water.  Filter  off  the  small 
amount  of  undissolved  mercuric  iodide  and  allow  the  double  salt 
to  crystallize  in  a  vacuum  desiccator  over  sulphuric  acid,  breaking 
up  occasionally  the  crust  which  forms  over  the  surface.  When  a 
thick,  pasty  mass  of  crystals  is  obtained,  drain  off  the  liquid  with 
suction,  and  without  washing  dry  the  light-yellow,  prismatic 
crystals  in  a  desiccator.  Obtain  more  of  the  product  from  the 
mother-liquor. 

Test-Tube  Experiments.  1.  Treat  a  little  mercuric  oxide  with  a 
few  cubic  centimeters  of  potassium  iodide  solution.  A  colorless, 
strongly  alkaline  solution  containing  potassium  mercuric  iodide 
and  potassium  hydroxide  is  produced.  The  formation  of  the  com- 
plex salt  makes  it  possible  for  the  mercuric  oxide  to  dissolve  with 
the  liberation  of  free  potassium  hydroxide  (cf.  No.  112). 

2.  Heat  a  sample  of  the  dry  salt  in  a  test-tube.  In  addition  to 
a  little  water,  mercuric  iodide  distils  off  and  condenses  in  the  yellow 
modification  on  the  cooler  part  of  the  tube.  After  some  time,  or 
more  quickly  on  touching  it  with  a  glass  rod,  the  sublimate  changes 
into  the  red  form,  and  this  can  again  be  changed  into  the  yellow 
condition  by  heating  (cf.  Prep.  17). 

107.   Potassium  Cobalticyanide,  Ks[Co(CN)J. 

Triturate  30  g.  of  cobalt  carbonate  with  a  little  water  until  it  is 
thoroughly  wet;  then  suspend  it  in  100  c.c.  of  water  and  dissolve  it 
by  adding  a  solution  of  110  g.  potassium  cyanide  in  400  c.c.  of  water 
Oxidize  the  potassium  cobaltocyanide  thus  formed  by  drawing  a 
vigorous  current  of  air  through^ the  liquid  for  an  hour.  After  filter- 
ing the  dark-yellow  solution,  add  to  it  40  g.  of  glacial  acetic  acid 
and  evaporate  it  to  crystallization  under  a  well- ventilated  hood 
outside  of  the  general  laboratory.  Drain  the  crystals,  and  wash 
them  with  alcohol  of  about  66%  by  volume.  Work  up  the  mother- 
liquor  repeatedly  as  long  as  a  sample  of  the  crystals  obtained  gives 
a  deep-blue  solution  when  heated  with  concentrated  sulphuric  acid. 
The  easily  soluble  potassium  acetate  remains  in  the  residual  liquor. 
Recrystallize  the  combined  fractions  from  a  solution  which  is 
slightly  acidified  with  acetic  acid.  Yield,  about  60  g.  The  product 
may  be  used  in  No.  109. 


146  COMPLEX  HALOGEN  ACIDS  AND  SALTS. 

108.   Hydroferrocyanic  Acid,  H4[Fe(CN)J. 

Hydroferrocyanic  acid  can  be  obtained  pure  without  difficulty,  and  dif- 
fers in  this  respect  from  most  of  the  related  acids  which  are  only  stable  in 
the  form  of  their  salts.  Hydroferrocyanic  acid  is  easily  soluble  in  water 
and  alcohol;  but  it  can  be  readily  precipitated  as  an  addition  product  with 
ether,  and  the  ether  can  be  removed  from  the  compound  by  heating  it  in  a 
current  of  hydrogen  at  80-90°. 

Treat  a  solution  containing  42  g.  of  potassium  ferrocyanide 
(0.1  mol.)  in  350  c.c.  of  water  with  120  c.c.  of  concentrated  hydro- 
chloric acid;  if  any  precipitate  of  potassium  chloride  separates, 
redissolve  it  by  adding  a  little  more  water.  Cool  and  add  about 
50  c.c.  of  ether.  After  standing  for  several  hours,  colorless,  glisten- 
ing, microscopic  crystals  separate  which  should  be  drawn  off  and 
washed  with  dilute  hydrochloric  acid  containing  a  little  ether.  To 
remove  any  admixed  potassium  chloride,  dissolve  the  product  in 
50  c.c.  of  alcohol,  filter,  precipitate  again  with  50  c.c.  of  ether, 
drain  off  the  liquid,  and  wash  the  crystals  with  the  ether.  Bring  the 
ether-hydroferrocyanic-acid  compound  into  an  Erlenmeyer  flask, 
which  is  provided  with  an  inlet  and  an  outlet  tube,  and  heat  it  in  an 
atmosphere  of  dry  hydrogen  at  80°  to  90°  on  the  water  bath.  The 
ether  is  removed  in  this  way  in  about  an  hour.  Yield,  about  12  g. 
The  hydroferrocyanic  acid  is  nearly  colorless  at  first,  but  it  quickly 
develops  a  light-blue  color  by  contact  with  the  air. 

109.   Addition  Products  of  Complex  Hydro-metal-cyanic  Acids 
with  Oxygen  Compounds. 

Some  of  the  hydro-metal-cyanic  acids  (hydroferrocyanic,  hydroferricyanic, 
and  hydrocobalticyanic  acids)  give  with  organic  oxygen  compounds  the 
most  varied  classes  of  more  or  less  stable  precipitates.1  It  has  been  assumed 
that  oxygen  in  these  compounds  is  tetravalent,  yet  this  is  by  no  means 
proved.  Some  of  these  compounds  have  been  analyzed,  for  instance  that 
between  one  molecule  of  hydroferrocyanic  acid  and  three  molecules  of  fur- 
furol.2  The  result  of  recent  investigations  indicates  that  such  precipitates 
consist  in  part  of  solid  solutions,  in  part  of  crystalline  double  compounds.3 

1.  Addition  Products  of  Hydroferrocyanic  Acid.  Dissolve  2.5  g. 
of  potassium  ferrocyanide  in  17  c.c.  of  water,  add  to  the  cold  solu- 
tion 8  g.  of  concentrated  hydrochloric  acid,  and  filter. 

1  v.  Baeyer  and  Villiger,  Ber.  34,  2687  (1901). 

2  Wagener  and  Tollens,  Ber.  39,  413  (1906). 

3  Mclntosh,  J.  Am.  Chem.  Soc.  30,  1097  (1908). 


POTASSIUM    COBALTOTHIOCYANATE.  147 

A  little  of  this  solution  when  treated  with  ether  gives  thin 
colorless  tablets  (compare  the  preceding  preparation);  with  ben- 
zaldehycle  and  concentrated  hydrochloric  acid  an  amorphous 
precipitate  is  formed  (microscope). 

2.  Addition    Products    of    H ydroferricyanic    Acid.     Prepare    a 
solution  of  the  free  acid  as  above  from  2.5  g.  of  potassium  ferri- 
cyanide,  6  c.c.  of  water,  and  7.5  g.  of  concentrated  hydrochloric 
acid.     A  little  of  this  solution  gives  with  ether  a  brown,  oily  mass; 
with  amyl  alcohol  and  a  large  amount  of  concentrated  hydro- 
chloric acid,  dull-yellow  prisms;  with  benzaldehyde  and  hydrochlo- 
ric acid,  small  tablets  which  quickly  become  greenish  yellow. 

3.  Addition   Products  of  Hydrocobalticyanic  Acid.     Prepare  a 
solution  of  the  free  acid  as  above  from  3  g.  of  potassium  cobalti- 
cyanide  (No.   107),  9  c.c.  of  water,  and   15  g.  of  concentrated 
hydrochloric  acid. 

A  little  of  this  solution  gives  with  ether,  flat  colorless  needles; 
with  benzaldehyde,  short  prisms.  By  shaking  equal  volumes  of 
the  hydrocobalticyanic  acid  solution  and  a  solution  of  camphor  in 
benzene  for  two  hours,  crystals  are  obtained  of  the  composition 
2  C]0H160  .  H3[Co(CN)6] .  2  H20;  the  crystals  may  be  washed  with 
benzene  and  20%  hydrochloric  acid. 

110.    Cobaltous  Salt  of  Hydromereurithiocyanic  Acid, 

Co[Hg(SCN)J. 

Prepare  two  solutions,  one  with  30  g.  of  mercuric  chloride  and 
44.5  g.  of  potassium  thiocyanate  in  500  c.c.  of  water,  and  another 
with  20  g.  of  cobalt  nitrate  in  50  c.c.  of  water;  then  mix  the 
two  clear  solutions;  a  shower  of  small,  deep-blue  crystals  of  the 
difficultly  soluble  double  salt  begins  to  fall  at  once.  After  standing 
12  hours,  drain  the  crystals,  wash  them  with  water,  then  with 
alcohol,  and  dry  the  product  in  the  steam  closet.  The  yield  is 
nearly  quantitative. 

111.   Potassium  Cobaltothiocyanate,  K2[Co(SCN)J. 

Prepare  a  hot,  saturated  solution  of  14.5  g.  of  crystallized  cobalt 
nitrate  and  24  g.  of  potassium  thiocyanate;  let  it  cool,  and  after 
several  hours  remove  the  crystals  of  potassium  nitrate  which  have 
separated  and  rinse  them  with  40  g.  of  amyl  acetate,  added  hi  small 
portions,  until  they  are  nearly  colorless.  Add  these  washings  to 


148  COMPLEX  HALOGEN  ACIDS  AND   SALTS. 

the  main  part  of  the  liquid  in  a  separatory  funnel  and  shake  thor- 
oughly. Separate  the  two  layers  that  form;  this  will  require  very 
close  attention,  since  both  are  deep  blue.  Extract  the  aqueous 
layer  once  or  twice  more  with  10  to  15  c.c.  of  amyl  acetate.  To 
obtain  a  good  product,  it  is  necessary  before  evaporating  the  amyl 
acetate  solution,  to  free  it  mechanically  from  drops  of  aqueous 
solution  by  pouring  it  back  and  forth  into  clean,  dry  beakers.  In 
order  to  crystallize  the  salt,  evaporate  off  one-tenth  of  the  amyl 
acetate  under  the  hood,  and  to  the  cooled  solution  add  slowly 
50-60  c.c.  of  low-boiling  ligroin.  Drain  off  the  crystals,  wash 
them  with  ligroin,  and  dry  them  over  sulphuric  acid.  The  product 
consists  of  deep-blue  needles;  when  dissolved  in  water,  it  disso- 
ciates and  a  red  solution  is  formed;  upon  the  addition  of  potassium 
thiocyanate  the  dissociation  is  driven  back  and  the  solution 
becomes  blue. 

112.    Cadmium  Iodide,  Cd[CdI4]  (Autocomplex  Compound). 

Cadmium  iodide  is,  according  to  Hittorf,  the  cadmium  salt  of  the  com- 
plex hydro-cadmi-iodic  acid,  H2[CdIJ.  A  compound  of  this  nature,  in  which 
one  and  the  same  metal  exercises  different  functions,  is  known  as  auto- 
complex.  (W.  Biltz,  1902.)  In  harmony  with  the  complex  nature  of  cadmium 
iodide,  the  simple  ions  Cd++  and  I~show  a  strong  tendency  to  combine  with 
each  other,  as  is  illustrated,  for  example,  by^the  fact  that  when  a  solution  of 
potassium  iodide  is  treated  with  cadmium  hydroxide,  the  liquid  becomes 
strongly  alkaline,  due  to  the  formation  of  undissociated  cadmium  iodide 
and  potassium  hydroxide. 

Allow  several  rods  of  pure  zinc  to  stand  for  about  24  hours  in 
contact  with  a  solution  of  26  g.  crystallized  cadmium  sulphate  in 
100  c.c.  of  water,  until  all  the  cadmium  is  precipitated  as  spongy 
metal.  To  test  for  complete  precipitation  treat  a  few  drops  of  the 
solution  with  hydrogen  sulphide;  no  cadmium  sulphide  should 
form.  Purify  the  finely  divided  cadmium  by  boiling  it  repeatedly 
with  water. 

Boil  the  cadmium  with  24  g.  of  iodine  and  50  c.c.  of  water,  in  a 
flask  with  return  condenser,  until  all  the  metal  is  dissolved  (1  or 
2  hours).  Then  continue  the  boiling  in  an  open  flask  to  remove 
any  excess  of  iodine,  filter,  and  concentrate  the  filtrate  to  crys- 
tallization. Work  up  the  mother-liquor.  Yield,  30  to  35  g.  of 
nearly  colorless,  lustrous  plates. 

To  show  that  the  amount  of  cadmium  ions  in  a  concentrated 


SODIUM   COBALTINITRITE.  149 

cadmium  iodide  solution  is  very  small,  treat  the  last  mother- 
liquor  with  hydrogen  sulphide.  A  slight  separation  of  cadmium 
sulphide  takes  place  after  some  time,  but  the  precipitation  is 
incomplete. 

NITRITO  ACIDS  AND  THEIR  SALTS. 

113.   Potassium  Mercurinitrite,  K3[Hg(NO2)5].  H2O. 

Treat  43  g.  of  yellow  mercuric  oxide  in  a  500  c.c.  flask  with  71  g. 
of  potassium  nitrite  and  240  g.  of  10%  acetic  acid,  and  shake 
frequently  until  all  is  dissolved.  Filter  and  concentrate  the 
solution  to  crystallization  on  the  water  bath.  The  separation  of 
crystals  is  much  aided  by  the  cautious  addition  of  some  alcohol. 
Collect  the  crystals  and  work  up  the  mother-liquor.  Recrys- 
tallize  the  combined  crude  product  from  a  little  water  to  which  a 
few  drops  of  potassium  nitrite  solution  have  been  added.  Pale- 
yellow,  lustrous  prisms,  or  plates,  are  obtained  which  are  easily 
soluble  in  water;  by  slow  evaporation,  crystal  aggregates  are 
sometimes  obtained  of  finger  length. 

The  small  content  of  mercuric  ions  in  the  solution  of  this  salt  is 
shown  by  its  indifference  toward  a  solution  of  urea,  or  toward  a 
cold  sodium  bicarbonate  solution  that  is  saturated  with  carbon 
dioxide;  on  the  other  hand,  precipitates  are  produced  by  sodium 
carbonate  and  by  sodium  hydroxide.  From  this  it  follows  that 
a  potassium  mercurinitrite  solution  contains  mercuric  ions  in 
about  the  same  concentration  as  a  solution  of  mercuric  chloride. 

114.   Sodium  Cobaltinitrite ;   Potassium  Cobaltinitrite. 

Sodium  Cobaltinitrite.  Dissolve  50  g.  of  cobaltous  nitrate  and 
150  g.  of  sodium  nitrite  in  150  c.c.  of  water.  Cool  the  solution 
to  40°  and,  while  shaking  frequently,  add  50  c.c.  of  50%  acetic 
acid,  a  little  at  a  time.  Then  oxidize  the  cobaltous  salt  by  draw- 
ing air  through  the  liquid  for  half  an  hour  (see  No.  107).  After 
some  time  filter  off  the  precipitate,  which  consists  of  sodium  cobalti- 
nitrite  and  perhaps  a  little  of  the  corresponding  potassium  salt. 
Stir  this  precipitate  with  50  c.c.  of  water  at  70°-80°  and  after 
10  minutes  filter  off  the  solution.  Combine  the  two  filtrates,  and 
by  introducing  350  c.c.  of  96%  alcohol  from  the  jet  of  the  wash- 
bottle  while  stirring,  throw  down  the  sodium  Cobaltinitrite  from 


150  NITRITO  ACIDS  AND   SALTS. 

the  solution.  After  standing  2  hours  collect  the  precipitate  and 
wash  it  twice  with  25  c.c.  of  alcohol.  Yield,  51  to  52  g. 

Purify  the  crude  product  by  recrystallizing  it  in  three  portions, 
each  of  which  is  stirred  up  with  1.5  times  its  weight  of  cold  water. 
Filter  off  the  small  undissolved  residues,  an,d  precipitate  each  of 
the  clear  filtrates  by  injecting  from  the  wash-bottle,  as  above, 
50  c.c.  of  a  mixture  of  alcohol  and  a  little  glacial  acetic  acid 
(70  :  1).  Wash  the  precipitate  with  alcohol  and  ether,  and  dry 
it  at  a  temperature  not  exceeding  80°.  Yield,  about  40  g. 

Potassium  Cobaltinitrite.  To  the  combined  mother-liquors  from 
the  above,  add  potassium  chloride  solution  until  precipitation  is 
complete.  Collect,  wash,  and  dry  the  yellow,  crystalline  precipitate. 

The  aqueous  solution  of  the  readily  soluble  sodium  cobaltinitrite 
is  a  useful  reagent  in  testing  for  the  presence  of  potassium  ions. 

115.   Potassium  Tetranitrito-diammine-cobaltate, 

K[Co(N02)4(NH3)2]. 

Dissolve*  10  g.  of  cobaltous  carbonate  in  a  barely  sufficient 
amount  of  hydrochloric  acid,  so  that  a  trace  of  residue  remains 
undissolved  and  the  solution  is  only  faintly  acid.  Dilute  the 
solution  to  200  c.c.  with  water  and  add  70  g.  of  ammonium  chloride, 
whereupon  the  color  changes  from  red  to  violet.  Warm  the 
liquid  to  50°  and  add  a  solution  of  100  g.  of  potassium  nitrite  in 
100  c.c.  of  water  which  is  likewise  warmed  to  50°  (use  a  large 
beaker  on  account  of  foaming).  Maintain  the  mixture  at  50°  for 
half  an  hour  and  then  place  it  in  the  ice-chest  for  24  hours.  Dull- 
brown  crystals  and  a  fine  yellow  powder  separate;  the  latter  may 
be  removed  by  rotating  the  mass  two  or  three  times  with  100  c.c. 
of  cold  water  and  each  time  pouring  the  liquid  and  the  suspended 
matter  away  from  the  brown  crystals. 

Dissolve  the  crude  product  in  150  c.c.  of  boiling  water;  filter 
and  cool  the  filtrate  at  once,  because  the  salt  is  decomposed  in  hot 
solution.  The  main  part  of  the  salt  crystallizes  on  cooling;  wash 
it  with  a  little  water  and  then  with  alcohol.  The  addition  of 
alcohol  and  a  little  ether  to  the  mother-liquor  causes  the  remainder 
of  the  salt  to  precipitate,  and  this  should  be  recrystallized  from  a 
little  water.  Yield,  8  g.  of  brown,  lustrous  crystals. 

A  cold  saturated  aqueous  solution  gives  yellowish-brown  prisms 
or  rhombic  leaflets  when  treated  with  mercurous  nitrate  solution. 


SILICOTUNGSTIC  ACID.  151 

CONDENSED  ACIDS  AND  THEIR  SALTS. 

Among  the  compounds  with  complex  anions  belong  a  great  number  of 
complicated  acids,  many  of  them  known  only  in  the  form  of  their  salts. 
The  acids  may  be  imagined  as  produced  by  the  condensation  of  oxy-acids 
with  their  own  anhydrides  or  with  the  anhydrides  of  other  oxy-acids  in  the 
most  varied  proportions;  the  components  which  come  chiefly  under  consid- 
eration in  this  connection  are  silicic,  stannic,  phosphoric,  vanadic,  molybdic, 
and  tungstic  acids,  and  their  anhydrides.  Pyrosulphuric  and  pyrochromic 
acids  may  be  regarded  as  the  simplest  type  of  the  condensed  acids,  but 
with  the  reservation  that  the  mechanism  of  the  reaction  assumed  for  their 
formation  (separation  of  water  from  two  molecules  of  the  simple  acid)  does 
not  of  necessity  hold  for  the  other  acids.  It  is  better,  therefore,  to  refrain 
from  attempting  to  ascribe  rational  formulas  to  such  substances,  and  merely 
to  give  their  constituents  one  after  the  other,  as,  for  example,  silicotungstic 
acid,  4  H2SiO3 .  12  WO3 .  29  H2O.  Of  the  more  complicated  compounds  of 
this  series,  phosphomolybdic  acid  is  the  best  known.  A  characteristic 
of  many  of  the  acids  of  this  class  is  the  ability  (see  discussion  under  No. 
109)  to  combine  with  ether.  Use  is  made  of  this  property  in  preparing 
the  pure  acids  (cf.  No.  117).  In  organic  chemistry  some  of  the  con- 
densed acids  are  used  in  precipitating  bases  of  high  molecular  weight ;  thus 
phospho-molybdic  and  phospho-tungstic  acids  serve  in  the  detection  of 
alkaloids.  The  latter  is  also  used  for  separating  the  decomposition  products 
of  proteins,  whereby  the  diamino-acids  give  more  difficultly  soluble  pre- 
cipitates than  the  monoamino-acids. 

116.   Ammonium  Phosphomolybdate,  (NH4)3PO4 . 12  MoO3 .  H2O. 

Treat  a  concentrated  solution  of  24  g.  ammonium  molybdate  with 
concentrated  nitric  acid  until  the  precipitate  of  molybdic  acid 
first  appearing  has  redissolved  and  the  solution  remains  clear  even 
on  boiling.  On  adding  a  solution  of  3.6  g.  of  sodium  phosphate  the 
well-known,  yellow  precipitate  is  produced.  Wash  this  and  dry 
it  in  the  steam  closet.  Yield,  about  20  g. 

117.    Silicotungstic  Acid,  4  H2SiO3 . 12  WO3 . 29  H2O. 

First  prepare  silicic  acid  hydrogel  by  dissolving  50  g.  of  com- 
mercial, precipitated  silicic  acid  in  12  g.  of  sodium  hydroxide  and 
50  c.c.  of  water,  diluting  to  200  c.c.,  heating  to  about  100°,  and 
precipitating  with  hydrochloric  acid.  Dilute  further  with  hot 
water,  drain  the  hydrogel  on  linen  cloth  in  a  large  Biichner  funnel 
and  wash  with  hot  water.  Stir  up  the  hydrogel  in  a  dish  with 
hot  water  and  again  drain  it  on  the  linen  cloth. 

Neutralize  a  solution  of  50  g.  pure,  commercial  sodium  tungstate 


152          ORGANOCOMPLEX  COMPOUNDS. 

in  200  c.c.  of  water  with  about  80  c.c.  of  2-normal  sulphuric  acid, 
using  litmus  paper  as  indicator.  Have  this  solution  boiling,  add 
the  hydro  gel  above  obtained,  and  boil  the  mixture  until  a  filtered 
sample  no  longer  gives  a  yellow  precipitate  of  tungstic  acid  when 
acidified  with  hydrochloric  acid.  During  the  boiling  add  a  little 
2-normal  sulphuric  acid  whenever  the  solution  becomes  alkaline 
(in  all  about  20  c.c.).  This  operation  should  take  half  to  three- 
quarters  of  an  hour.  Water  should  be  added  from  time  to  time 
to  replace  that  lost  by  evaporation.  Finally,  filter  the  solution. 

Cool  the  filtrate,  which  amounts  to  about  200  c.c.  in  volume,  and 
acidify  it  strongly  by  adding  10  to  20  c.c.  of  concentrated  sulphuric 
acid.  Again  cool  and  add  ether,  little  by  little  with  constant 
shaking,  until  on  settling  in  a  separatory  funnel  three  layers  are 
formed:  on  top  an  ethereal,  in  the  middle  an  aqueous  solution,  and 
underneath  a  thick,  oily  layer  of  a  liquid  compound  of  silico- 
tungstic  acid  with  ether  (cf.  No.  109).  If  this  lowest  layer  does  not 
form,  or  forms  but  incompletely,  more  sulphuric  acid  must  be  added. 
Allow  the  whole  to  stand  over  night,  then  draw  off  the  bottom 
layer  into  a  dry  beaker,  free  it  from  drops  of  water  by  pouring  it 
back  and  forth  into  dry  beakers,  and  remove  the  ether  by  warming 
it  for  some  time  upon  the  water  bath.  Dissolve  the  residue  in  a 
little  water,  filter  the  solution,  concentrate  it  to  a  small  volume, 
and  bring  it  to  crystallization  by  surrounding  the  beaker  with 
ice.  If  no  crystals  separate,  the  solution  must  be  concentrated 
still  further.  Drain  the  crystals  quickly  from  the  liquid;  do  not 
wash  this  product,  but  free  it  from  mother-liquor  by  spreading  it 
on  an  unglazed  plate.  After  the  crystals  have  become  white  on 
the  porous  plate,  dry  them  completely  by  letting  them  remain  in 
a  vacuum  desiccator  over  sulphuric  acid.  Yield,  13  to  14  g. 

When  working  with  larger  amounts  of  silicotungstic  acid,  large 
characteristic  crystals  can  be  obtained  by  slow  evaporation  in  a 
vacuum  desiccator. 

ORGANOCOMPLEX  COMPOUNDS. 
118.    Potassium  Ferric  Oxalate,  K3[Fe(C2O4)3];  Plantiotypes. 

Oxidize  a  solution  of  35  g.  of  crystallized  ferrous  sulphate  in 
100  c.c.  of  water  by  boiling  it  with  just  the  necessary  amount  of 
nitric  acid  (test  with  potassium  ferricyanide).  Dilute  the  liquid 
to  2  liters,  add  ammonia  and  wash  the  precipitate  of  hydrated  ferric 


OPTICAL  ROTATION   OF  URANYL  MALATE.  153 

oxide  by  decantation  for  several  days,  then  collect  it  on  a  large, 
plaited  filter  and  wash  it  with  hot  water. 

To  a  hot  solution  of  44  g.  of  crystallized  acid  potassium  oxalate 
in  100  c.c.  of  water  add  the  ferric  oxide  hydrogel  a  little  at  a  time 
until  no  more  will  dissolve.  Such  a  solution  is  sensitive  to  direct 
sunlight.  Filter  and  concentrate  the  nitrate  to  crystallization, 
wash  the  emerald-green  crystals  with  water  and  alcohol,  and  dry 
them  in  a  vacuum  desiccator  over  sulphuric  acid. 

Platinotypes. 

Potassium  ferric  oxalate  is  changed  by  the  action  of  light  to  potassium, 
ferrous  oxalate,  and  the  latter  reduces  platinum  salts  to  metallic  platinum. 

Soak  a  piece  of  filter  paper,  the  size  of  a  photographic  plate, 
with  potassium  ferric  oxalate  solution  and  dry  it  in  the  hot  closet. 
Then  place  the  paper  in  a  shallow  glass  tray;  wet  it  uniformly  with 
about  2  c.c.  of  a  5%  chloroplatinic  acid  solution,  whereupon  it 
turns  yellowish  red;  and  dry  it  again,  avoiding  any  strong  illumi- 
nation. Then  expose  it  for  about  an  hour  to  a  medium  light  under 
a  paper  stencil  in  a  printing  frame.  At  the  end  of  the  exposure, 
only  the  outlines  of  the  pattern  are  to  be  distinguished,  but  the 
print  may  be  developed  by  placing  it  for  about  one  minute  in  a 
warm  solution  of  potassium  oxalate.  Fix  the  "  picture  "  in  dilute 
hydrochloric  acid,  wash  it  repeatedly  with  water  and  then  dry  it. 
If  potassium  chloroplatlnite  is  used  instead  of  chloroplatinic  acid 
the  picture  is  developed  of  itself  during  the  printing,  but  it  should 
be  fixed  as  in  the  preceding  case  by  washing  with  dilute  hydro- 
chloric acid  and  water. 

119.    Optical  Rotation  of  TJranyl  Laevo-malate. 

Aqueous,  not  too  concentrated,  solutions  of  ordinary  malic  acid  rotate 
the  plane  of  polarized  light  feebly  toward  the  left.  The  presence  of  uranyl 
salts  increases  the  extent  of  the  rotation  very  considerably,  and  this  is  presum- 
ably due  to  the  formation  of  complex  compounds,  although  such  substances 
have  not  yet  been  isolated.  For  this  experiment  a  sensitive  polarizing  appa- 
ratus with  graduations  is  necessary. 

Prepare  the  following  solutions:  (I)  1.3  g.  of  malic  acid  in  10  c.c. 
of  water;  (II)  4  g.  of  uranyl  nitrate  in  10  c.c.  of  water;  (III)  2  g. 
of  potassium  hydroxide  in  20  c.c.  of  water. 

First  mix  1  c.c.  of  (I)  and  2  c.c.  of  (III)  and  dilute  to  20  c.c. 


154          ORGANOCOMPLEX  COMPOUNDS. 

Next  prepare  the  same  mixture  of  solutions  (I)  and  (III),  add 
1.5  c.c.  of  solution  (II)  and  dilute  to  20  c.c.  Determine  in  both 
samples  the  degree  of  rotation.  If  the  polariscope  is  not  very 
sensitive,  use  instead  of  the  first  mixture  one  which  contains  5  c.c. 
of  solution  (I). 

If  c  is  the  concentration  of  the  malic  acid,  I  the  length  of  the 
tube  (usually  20  cm.),  and  aD  the  angle  of  rotation  read  with 

sodium  light,  then  the  specific  rotation  is  [a]D  = —•  -      The 

experiment  with  malic  acid  gives  for  the  value  of  [aD]  about 
—  3°,  with  uranyl-malic  acid,  about  —  475°. 


CHAPTER  V. 


COMPOUNDS  CONTAINING  A  COMPLEX  POSITIVE 
COMPONENT. 

THE  fundamental  principles  involved  in  the  formation  and  dissociation  of 
compounds  containing  complex  cations  are  essentially  the  same  as  in  com- 
pounds with  complex  anions.  For  example,  hexamminenickelous  bromide 
is  formed  as  represented  by  the  equation : 

NiBr2  +  6  NH3  =  [Ni(NH,)  J  Br2; 

hexaaquochromic  chloride: 

CrCl3  +  6  H20  =  [Cr(H20)6]  C13; 

ammonium  chloride  is  formed  either  in  a  similar  manner: 
NH3  +  HC1  =  [NH3.H]C1; 

or  by  the  association  of  the  ions: 

NH4+  +  Cl-  =  NH4C1. 

The  last  two  equations,  viewed  in  the  light  of  Werner's  theory  of  secondary 
valences  (cf.  page  141),  lead  to  the  following  structural  formula: 


[~H\  1 

H— N-H-    -Cl. 
LH' 


Compounds  with  complex  cations  dissociate  according  to  reactions  which 
are  the  reverse  of  the  reactions  of  their  formation;  naturally  both  formation 
and  dissociation  may  take  place  in  stages. 

In  the  production  of  complex  cations,  ammonia  and  water  come  most 
frequently  into  consideration.  Among  the  ammonia  compounds,  those  in 
which  ammonia,  or  a  substituted  ammonia,  is  joined  to  hydrogen  (ammonium 
or  substituted  ammonium  compounds)  are  to  be  distinguished  from  those 
in  which  ammonia,  or  a  substituted  ammonia,  is  joined  to  metal.  The  classi- 
fication which  follows  is  based  upon  this  distinction. 

Other  complex-producing  substances  than  those  mentioned  are  less  often 
met  with.  As  representative  of  the  compounds  which  they  form,  the  nitric- 
oxide-metal  compounds  (cf.  the  well-known  test  for  nitric  acid  with  ferrous 
sulphate)  may  be  mentioned,  as  well  as  the  addition  products  with  alcohol, 
ether,  and  ethyl  acetate,  which  are  similar  in  nature  to  the  hydrates. 

155 


156  AMMONIUM  COMPOUNDS. 

AMMONIUM    COMPOUNDS    AND    SUBSTITUTED  AMMONIUM 

COMPOUNDS. 

120.  Dissociation  of  Ammonium  Chloride. 

The  fact  that  the  vapor  density  determination  of  ammonium  chloride 
shows  the  molecular  weight  to  be  but  one-half  the  formula  weight,  leads  to 
the  conclusion  that  the  number  of  molecules  is  doubled  by  the  dissociation  of 
the  substance  into  ammonia  and  hydrogen  chloride.  If  the  compound  is 
volatilized  into  an  atmosphere  either  of  ammonia  or  of  hydrogen  chloride,  the 
dissociation  is  driven  back  in  accordance  with  the  mass-action  law  —  this 
phenomenon  being  especially  pronounced  when  working  at  lower  temper- 
atures and  under  reduced  pressure.  In  the  complete  absence  of  water  the 
dissociation  fails  to  take  place,  thus  showing  in  a  remarkable  manner  the 
catalytic  effect  of  traces  of  moisture. 

The  dissociation  of  ammonium  chloride  may  be  demonstrated  qualita- 
tively, by  taking  advantage  of  the  greater  velocity  at  which  the  lighter 
ammonia  diffuses  as  compared  with  the  heavier  hydrogen  chloride;  the 
products  of  diffusion  may  be  most  conveniently  separated  by  the  use  of  a 
diaphragm. 

Fasten  a  piece  of  combustion  tubing,  20  cm.  long,  in"  a  hori- 
zontal position.  Insert  a  loose  plug  of  asbestos  at  the  middle  of 
the  tube,  and  place  about  a  gram  of  ammonium  chloride  on  one 
side  of  the  plug.  Heat  the  asbestos  diaphragm  and  the  ammonium 
chloride  by  means  of  a  wide  burner  so  that  a  slow  sublimation 
takes  place  and  the  entire  tube  becomes  filled  with  the  vapors. 
After  a  few  minutes  test  the  gases  at  both  ends  of  the  tube  with 
moist  litmus  paper.  An  acid  reaction  is  shown  on  the  side  of 
the  asbestos  plug  on  which  the  solid  salt  was  placed,  while  on  the 
other  side  an  alkaline  reaction  is  obtained. 

121.   Hydroxylamine  Sulphate,  [NH2OH.H]2SO4. 

Sodium  nitrite  and  sodium  bisulphite  react  together  in  cold  aqueous  solu- 
tion, at  first  molecule  for  molecule,  forming  nitrososulphonate  of  sodium: 

ONjONaTH:SO3Na  =  NaOH  +  ON-SO3Na. 

Then,  by  the  immediate  taking  up  of  a  second  molecule  of  sodium  bisul- 
phite, the  stable  sodium  salt  of  hydroxylaminedisulphonic  acid  results: 
'ON-SO3Na  +  NaHSO3  =  HO-N(SO3Na)2. 

If  the  solution  is  warm  a  third  molecule  of  sodium  bisulphite  reacts  and 
nitrilo-sulphonate  of  sodium  is  formed. 

When  heated  above  100°  with  water  the  sodium  salt  of  hydroxylaminedi- 
sulphonic acid  is  hydrolyzed  into  hydroxylamine  and  sodium  bisulphate: 
HO-N(SO3Na),  +  2  H20  =  HONH2  +  2  NaHS04, 


HYDROXYLAMINE  SULPHATE  157 


and  these  substances  interact  to  form  hydroxylamine  sulphate,  [HONH3]2SO4, 
and  neutral  sodium  sulphate.  The  presence  of  barium  chloride  favors  this 
reaction,  since  the  sulphate  ions  are  thereby  precipitated  as  fast  as  they  are 
formed.  The  filtrate  then  contains,  in  addition  to  free  hydrochloric  acid 
only  hydroxylamine  hydrochloride,  HONH2-HC1,  and  sodium  chloride,  and 
these  can  be  separated  by  means  of  alcohol. « 

Hydroxylamine  is  of  great  importance  in  organic  chemistry,  where  it  is 
used  both  as  a  reducing  agent  and  as  a  reagent  for  the  carbonyl  groups  of 
aldehydes  and  ketones: 

(CH3)2C:OTH2|NOH  =  (CH3)2C  :NOH  +  H2O. 

acetone  acetone-oxime 

Hydrazine  (No.  122)  and  semicarbazid  (No.  123)  show  a  similar  behavior: 
(CH3)2CO  +  H2N-NH2  =   (CH3)2C  :N-NH2    +  H2O 

hydrazine  acetone-hydrazone 

2  (CH3)2CO  +  H2N-NH2  =     (CH3)2C  :  N-N:  C(CH3)2  +  2  H2O 

acetone-azine 

(CH3)2CO  +  H2N-NH-CO-NH2  =   (CH3)2C  :  N.NH-CONH2+H2O 

semicarbazide      •  acetone-sernicarbazone 

Saturate  a  solution  of  143  g.  crystallized  sodium  carbonate  in 
100  c.c.  water  with  sulphur  dioxide  (from  150  g.  copper  and  600  g. 
concentrated  sulphuric  acid),  and  allow  the  resulting  solution  to 
drop  from  a  funnel  into  a  very  cold  solution  of  36  g.  sodium 
nitrite  in  60  g.  water;  keep  the  latter  solution  cooled  by  surround- 
ing the  beaker  with  a  mixture  of  salt  and  ice;  at  no  time  should 
the  temperature  of  the  liquid  rise  above  0°.  Stir  the  mixture 
vigorously  during  the  treatment,  best  by  means  of  a  mechanical 
stirrer.  At  the  last,  remove  any  excess  of  nitrite  by  conducting 
sulphur  dioxide  into  the  solution. 

Dilute  the  solution  to  one  liter,  place  it  in  a  round-bottopsed 
flask,  and  heat  it  to  boiling  upon  a  Babo  funnel.  To  the  boiling- 
liquid  add  a  hot  solution  of  barium  chloride  (about  250  g.  of  the 
crystallized  salt  in  300  c.c.  water),  avoiding  an  appreciable  excess, 
until  a  little  of  the  solution  on  being  filtered  is  found  free  from 
sulphate.  The  duration  of  this  operation  is  about  one  hour. 
Allow  the  liquid  to  settle,  filter  and  evaporate  the  filtrate,  at  first 
over  a  free  flame  and  finally  to  dry  ness  on  the  water  bath.  Break 
up  frequently  the  crusts  of  salt  which  are  formed,  and  draw  them 
up  on  to  the  sides  of  the  dish.  Place  the  anhydrous  residue  in  a 
flask  and  extract  it  three  times  by  boiling  it  with  alcohol,  using 
200,  100,  and  50  grams  respectively  of  the  latter;  evaporate  the 
alcoholic  filtrate  to  a  small  volume,  allow  the  solution  to  cool 


158  SUBSTITUTED    AMMONIUM  COMPOUNDS. 

completely,  collect  the  crystals  on  a  filter,  and  evaporate  the  cold 
mother-liquor  to  obtain  more  crystals.  Recrystallize  the  crude 
product  from  half  its  weight  of  hot  water.  Yield,  10  to  14  g.  of 
hydro xylamine  chloride. 

Heat  a  trace  of  the  hydroxylamine  salt  with  Fehling's  solution; 
cuprous  oxide  is  precipitated. 

Potassium  Salt  of  Hydroxylaminedisulphonic  Acid.  On  add- 
ing potassium  chloride  to  some  of  the  solution  of  the  sodium 
salt  of  hydroxylaminedisulphonic  acid  obtained  in  the  course  of 
the  foregoing  procedure,  small  crystals  of  the  potassium  salt 
HO  •  N:  (S03K)2  are  deposited. 

Acetoneoxime,  (CH3)2C:NOH.  Add  a  solution  of  8  g.  sodium 
hydroxide  in  12  c.c.  water  slowly  and  with  constant  stirring,  but 
without  cooling,  to  a  mixture  of  14  g.  hydroxylamine  hydrochloride, 
12  g.  acetone  and  13  c.c.  water.  Acetoneoxime  separates  out  as 
an  upper  layer  which  solidifies  after  standing  for  some  time. 
After  several  hours  pour  off  the  aqueous  solution  and  dissolve  the 
oxime  in  ether,  in  which  it  is  extremely  soluble.  Pass  this  solu- 
tion through  a  dry  filter  to  remove  any  suspended  drops  of  aqueous 
solution,  and  evaporate  the  ether  on  the  water  bath,  taking  the 
usual  precautions.  On  cooling,  the  oxime  solidifies  in  beautiful 
crystals;  if  desired,  it  may  be  recrystallized  from  a  low  boiling 
ligroin.  Yield,  10  to  13  g.  Melting-point  59°-60°. 

122.   Hydrazine  Sulphate,  [N2H4.H2]  SO4 ;  Monochloramine,  NH2  Cl. 

Hydrocyanic  acid,  as  formed  by  the  hydrolysis  of  potassium  cyanide, 
unites  in  aqueous  solution  with  two  molecules  of  potassium  bisulphite,  form- 
ing aminomethanedisulphonate  of  potassium: 

HCN  +  2  HSO3K  =  H(NH2)C(SO3K)2. 

By  strongly  acidifying  the  solution  with  hydrochloric  acid,  the  difficultly- 
soluble  acid  salt  is  precipitated : 

H(NH2)C(SO3H)SO3K. 

By  "diazotizing"  the  amido  group,  diazomethanedisulphonate  of  potassium  is 
obtained : 

N\ 

H(NH2)C(S03H)S03K  +  KNO2    =  =    ||   ,  C(SO3K)2  +  2  H2O. 

N 

This  salt  is  capable  of  adding  on  one  molecule  of  potassium  sulphite  at  the 
point  of  the  nitrogen  double  bond,  and  the  acid  corresponding  to  the  salt 


HYDRAZINE   SULPHATE.  159 

thus  formed  breaks  down  with  water,  essentially  into  sulphur  dioxide,  carbon 
dioxide,  and  hydrazine  sulphate: 

H.N 

!     .  C(S03H)2  +  H20  =  (H2N.NH2)H2S04  +  CO2  +  2  SO2. 

HSO3.N  '  hydrazine  sulphate 

Saturate  a  solution  of  75  g.  potassium  hydroxide  in  300  c.c.  of 
wrater  with  sulphur  dioxide.  To  the  solution  of  potassium  bisul- 
phite thus  prepared  add  50  g.  of  powdered  potassium  cyanide 
(98-99%)  while  shaking;  the  yellow  color  of  the  bisulphite  sol- 
ution disappears,  and  the  mixture  becomes  somewhat  heated. 
After  all  the  cyanide  has  dissolved,  heat  the  solution  on  the  water 
bath,  and  when,  after  some  time,  the  liquid  has  become  alkaline, 
acidify  it  cautiously  with  hydrochloric  acid.  Repeat  the  cautious 
additions  of  hydrochloric  acid  until  the  solution  has  become 
permanently  acid.  This  operation  requires  from  1.5  to  2  hours, 
and  in  all  about  30  c.c.  of  concentrated  hydrochloric  acid  are 
necessary.  Finally,  add  an  additional  150  c.c.  of  concentrated 
hydrochloric  acid  and  allow  the  mixture  to  stand  in  the  ice-chest, 
when  60  to  80  g.  of  the  crystals  of  acid  potassium  aminomethane- 
disulphonate  are  obtained.  Collect  this. product  on  a  filter  and 
wash  it  with  water. 

Treat  this  salt  in  separate  portions  of  23  g.  each  as  follows: 
stir  each  portion  to  a  paste  with  34  c.c.  of  water  and  then  add  to 
it  a  solution  of  10  g.  potassium  nitrite  in  6  c.c.  water.  The  tem- 
perature of  the  mixture  rises  slowly  to  40°-50°  and  within  10  or 
15  minutes  all  of  the  solid  salt  has  passed  into  solution.  Make  the 
solution  alkaline  with  a  little  caustic  potash  and  allow  it  to  cool; 
about  18  g.  of  potassium  diazomethanedisulphonate  are  deposited 
from  each  portion  in  the  form  of  orange-yellow  needles. 

Dissolve  the  latter  salt  in  a  solution  of  an  equal  weight  of 
crystallized  sodium  sulphite  in  two-thirds  as  much  water;  make 
the  solution  alkaline  by  the  addition  of  a  little  sodium  carbonate 
solution  and  then  warm  slightly  until  the  color  has  disappeared. 
Finally,  decompose  the  salt  of  the  trisulphonic  acid  by  adding  an 
amount  of  20%  sulphuric  acid  equal  to  five  times  the  weight  of 
the  salt.  When  the  liquid  no  longer  smells  of  sulphur  dioxide, 
filter  it  and  allow  it  to  cool,  whereby  the  hydrazine  sulphate 
crystallizes  out.  Obtain  a  further  yield  from  the  mother-liquor 
and  recrystallize  the  entire  crude  product  from  water.  Yield, 


160  SUBSTITUTED   AMMONIUM  COMPOUNDS. 

about  40%  of  the  weight  of  the  potassium  diazomethanedisul- 
phonate.     Dependent  preparation:    Sodium  Hydrazoate,  No.  72. 

Monochloramine,  NH2Cl;  Hydrazine  Sulphate  according  to 
Raschig.1 

Ammonia  is  chlorinated  by  the  action  of  sodium  hypochlorite  in  dilute 
aqueous  solution: 

NH3  +  NaOCl  =  NH2C1  +  NaOH. 

The  monochloramine  thus  formed  reacts  with  more  ammonia  to  form  hydra- 
zine  chloride: 

NH2C1  +  NH3  =  H2N  .  NH2 .  HC1. 

Place  600  g.  of  ice  and  a  cold  solution  of  85  g.  of  sodium  hy- 
droxide in  160  c.c.  of  water  in  a  one-liter  flask,  weigh  the  flask  and 
contents,  wrap  the  flask  in  a  towel,  and  pass  in  a  rapid  current  of 
chlorine  until  the  gain  in  weight  is  exactly  71  grams;  avoid  an 
excess  of  chlorine.  Then  on  diluting  to  one  liter,  an  approximately 
molal  solution  of  sodium  hypochlorite  is  obtained. 

Monochloramine.  Add  50  c.c.  of  this  sodium  hypochlorite 
solution  to  a  mixture  of  25  c.c.  2-normal  ammonia  and  75  c.c.  of 
water.  The  liquid  ceases  to  smell  of  ammonia,  and  in  its  place  a 
peculiar,  penetrating  odor  of  monochloramine  is  noticeable  which 
is  similar  to  that  of  nitrogen  chloride.  Nitrogen  is  evolved  from 
the  solution  at  the  same  time,  owing  to  the  fact  that  monochlora- 
mine decomposes  slowly  in  an  aqueous  solution  according  to  the 
following  equation : 

3  NH2C1  =  N2  +  NH4C1  +  2  HC1. 

Hydrazine  Sulphate.  Mix  200  c.c.  of  20%  ammonia,  5  c.c.  of 
a  1%  gelatin  solution,  and  100  c.c.  of  the  above  prepared  molal 
sodium  hypochlorite  in  a  one-liter  Erlenmeyer  flask,  heat  the 
mixture  immediately  to  boiling,  and  boil  it  for  half  an  hour.  The 
solution  is  thereby  concentrated  to  about  one-half  its  original 
volume.  Cool  it,  then  add  20  c.c.  of  2-normal  sulphuric  acid,  and 
surround  the  flask  with  ice.  Hydrazine  sulphate  crystallizes  out 
in  the  form  of  glistening,  transparent  crystals  which  are  obtained 
pure  by  recrystallizing  from  water.  Yield,  5  to  6  grams. 

This  process,  which  is  very  interesting  from  a  theoretical  standpoint,  is 
now  used  technically  for  the  production  of  hydrazine  sulphate,  whereby  the 
cost  of  the  latter  has  been  considerably  reduced. 

1  Raschig,  Ber.  40,  4586  (1907). 


SEMI  GARB  AZIDE  HYDROCHLORIDE  161 


123.   Semicarbazide  Hydrochloride. 

From  ammonia  and  cyanic  acid  urea  can  be  obtained,  with  the  interme- 
diate formation  of  ammonium  cyanate.  From  hydrazine,  which  is  an  amino 
substituted  ammonia,  semicarbazide 

/NH2 
C=  O 


an  analogous  substituted  urea  is  obtained.  Semicarbazide,  like  free  hydra- 
zine, condenses  with  carbonyl  groups  (cf.  Introduction  to  No.  121),  and  the 
semicarbazones  thus  formed  being  difficultly  soluble,  serve  for  the  detection 
and  isolation  of  carbonyl-containing  compounds. 

Dissolve  130  g.  hydrazine  sulphate  and  54  g.  anhydrous  sodium 
carbonate  in  50  c.c.  of  water,  warm  the  solution  to  50°—  60°,  and 
then  add  a  solution  of  86  g.  potassium  cyanate  in  500  c.c.  water. 
The  next  day  filter  off  the  few  grams  of  secondary  product  wrhich 
have  separated,  treat  the  filtrate  with  120  g.  of  acetone  and  allow 
the  mixture  to  stand  in  a  flask,  with  frequent  shaking,  for  another 
24  hours.  Drain  the  precipitated  salt,  and  evaporate  the  mother- 
liquor  to  dryness  on  the  water-bath,  stirring  towards  the  end. 
Place  the  entire  amount  of  the  salt  in  an  automatic  extraction 
apparatus,  and  extract  it  thoroughly  with  alcohol,  in  which  the 
acetone-semicarbazone  is  more  soluble  than  in  acetone;  a  few 
cubic  centimeters  of  acetone  should  be  mixed  with  the  alcohol. 
The  acetone-semicarbazone  crystallizes  in  the  distilling  flask  of 
the  extraction  apparatus.  After  it  is  drained  and  washed  with  a 
little  alcohol  and  ether,  it  melts  at  186°-187°.  Crystallize  the 
remainder  of  the  salt  by  evaporating  the  alcoholic  mother-liquor 
•somewhat  and  adding  a  little  ether.  Yield,  80%. 

The  above  product  can  be  transformed  quantitatively  into 
semicarbazide  hydrochloride,  as  follows:  Warm  the  acetone- 
semicarbazone  gently  with  concentrated  hydrochloric  acid  in  the 
proportion  of  11.5  g.  of  the  former  to  10  g.  of  the  latter  until  the 
solid  is  just  dissolved.  On  cooling,  the  solution  solidifies  to  a 
thick  paste  of  colorless,  well-formed  needles.  Drain  the  product 
with  suction  and  wash  it  with  a  little  alcohol  and  ether;  it  melts  at 
173°  with  decomposition.  To  the  mother-liquor,  add  double  its 
volume  of  alcohol,  and  bring  the  rest  of  the  salt  to  crystallization 
by  adding  ether. 


162  SUBSTITUTED  AMMONIUM  COMPOUNDS. 


124.   Millon's  Base. 

Millon's  base,  Hg2NOH  .  2  H2O,  is  formed  by  the  action  of  mercuric  oxide 
upon  an  ammonia  solution.  Since,  by  merely  heating,  this  compound  is 
changed  first  to  the  monohydrate  and  then  to  dimercuriammonium  hydrox- 
ide, Hg2NOH,  it  follows  that  in  Millon's  base  itself  the  mercury  is  probably 
united  to  the  nitrogen  atom,  as  is  unquestionably  the  case  in  the  dehydration 
product.  Whether  the  water  is  present  in  the  compound  as  water  of  crystalli- 
zation, or  whether  it  is  combined  by  atomic  valences,  has  not  been  established. 
In  the  latter  case  Millon's  base  would  be  regarded  as  dihydroxydimercuri- 
ammonium  hydroxide  (I)  and  its  first  dehydration  product  as  oxydimercuri- 
ammonium  hydroxide  (II)  ; 

HOHg\  /Hgx 

' 


HOHg 


NH2OH     (I)  O  ^NH2OH     (II). 


The  ability  of  mercury  to  replace  hydrogen  in  certain  compounds,  with  the 
formation  of  substances  like  the  above,  which  are  stable  in  presence  of  water, 
is  also  apparent  in  several  series  of  mercury  organic  compounds. 

Treat  a  solution  of  25  g.  mercuric  chloride  in  200  c.c.  of  water 
at  70°  with  a  solution  of  7.5  g.  sodium  hydroxide  in  20  c.c.  water, 
wash  the  resulting  precipitate  several  times  by  decantation, 
collect  it  on  a  suction  filter  and  wash  it  further  with  water.  Mean- 
while prepare  a  carbonate-free  solution  of  ammonia  by  distilling 
a  mixture  of  150  g.  concentrated  ammonia  with  20  g.  of  lime,  first 
placing  100  c.c.  of  cold  water  in  the  receiver.  Introduce  the 
moist  mercuric  oxide  into  this  ammonia  solution  and  allow  the 
mixture  to  stand  with  frequent  shaking  for  a  day  or  two  in 
the  dark.  Collect  the  product  on  a  suction  filter,  wash  it  with 
water,  alcohol  and  finally  with  ether,  and  dry  it  at  the  temperature 
of  the  laboratory. 

Salts  of  Millon's  Base. 

When  solutions  of  ordinary  salts  are  treated  with  Millon's  base,  salts  of  the 
latter,  which  are  very  difficultly  soluble,  and  free  metal  hydroxide,  are  formed. 

Shake  a  little  of  Millon's  base  with  a  dilute  solution  of  potassium 
iodide.  The  solution  becomes  alkaline,  and  the  brown  residue  is, 
in  all  probability,  identical  with  the  well-known  precipitate 
obtained  in  the  test  for  ammonia  with  Nessler's  reagent.  Filter 
the  solution  through  a  double  filter  of  hardened  paper,  and  clarify 
the  turbid  filtrate  by  shaking  it  with  pieces  of  torn  filter  paper  and 


METAL-AMMONIA  COMPOUNDS.  163 

again  filtering.  The  solution,  after  being  acidified  with  nitric 
acid,  gives  no  precipitate  with  silver  nitrate. 

A  dilute  solution  of  copper  sulphate,  or  one  containing  ferric 
chloride,  when  shaken  with  Millon's  base,  is  entirely  freed  from  all 
dissolved  salt. 

Oxydimercuriammonium  hydroxide.  Spread  10  g.  of  Millon's 
base  in  a  thin  layer  on  a  watch-glass  and  allow  it  to  dry  for  five 
days  over  lime  in  a  desiccator,  adding  in  the  first  place  0.5  to  1  c.c. 
of  concentrated  ammonia  to  the  lime  so  that  an  atmosphere  of 
ammonia  is  produced.  The  product  is  explosive.  Clamp  a  test- 
tube  containing  a  small  amount  of  the  substance  in  an  upright 
position  and  heat  it  behind  the  lowered  window  of  the  hood. 

METAL-AMMONIA    COMPOUNDS. 

The  importance  of  the  metal-ammonia  compounds  lies  in  their  great 
number,1  in  their  stability,  and  in  the  significance  of  the  theoretical  questions 
pertaining  to  them.  Among  the  best  known  of  these  compounds  are  those 
containing  trivalent  cobalt,  chromium,  rhodium  and  iridium.  Metal-ammonia 
compounds  are  formed  by  the  addition  of  gaseous  ammonia  to  the  solid  salts, 
or  by  the  union  of  the  two  components  in  solution. 

The  first  method,  inasmuch  as  it  represents  the  combination  of  two  sub- 
stances of  widely  different  volatility  (very  difficultly  volatile  metal  salt  and 
gaseous  ammonia) ,  offers  opportunity  for  the 
investigation  of  these  substances  from  the  NH3 
standpoint  of  heterogeneous  equilibrium.  • 
According  to  the  laws  of  heterogeneous  equi-  -. 
librium,  every  substance  which  gives  off  a 
gaseous  decomposition  product  has  at  a  given 
temperature  a  perfectly  definite  decomposition  3 
pressure  that  is  independent  of  the  amount  of 
the  decomposing  substance.  Or,  what  signifies  50°  160°  150°  260°  250°  300° 

the  same  thing,  it  possesses,  at  a  constant  Temperature 

pressure  (e.g.  atmospheric  pressure),  a  defi-  jrIG    22. 

nite  decomposition  temperature  which   may 

be  compared  with  the  boiling-point  of  a  liquid.  In  an  atmosphere  of  ammo- 
nia at  760  mm.  pressure,  the  compound  Zn(NH3)6Cl2  decomposes  at  59°, 
losing  two  molecules  of  NH3,  and  in  the  same  way  the  compound  Zn(NH3)4Cl. 
decomposes  at  89.5°  and  Zn(NH3)2Cl2  at  269°  (Fig.  22).  By  starting  with 
the  compound  richest  in  ammonia  and  heating  it  progressively  in  an  atmos- 
phere of  ammonia,  a  loss  in  weight  is  first  observed  to  take  place  at  59°,  which 
corresponds  to  the  escape  of  two  molecules  of  ammonia;  after  this  the  weight 

1  Werner  estimates  the  number  of  compounds  of  the  general  formula 
MXn(NH8)m  as  1700. 


164  METAL-AMMONIA  COMPOUNDS. 

remains  constant  until  the  second  decomposition  point,  89.5°,  is  reached.  On 
overstepping  this  point  a  third  constancy  of  weight  is  maintained  until  at 
269°  the  remaining  ammonia  molecules  escape  and  the  weight  above  tnis 
temperature  corresponds  to  that  of  the  ammonia-free  salt.  It  is  evident  that 
the  existence  of  ammonia  compounds  of  this  nature  may  be  discovered  by 
means  of  systematic  investigations  carried  out  in  this  manner. 

The  second  method  of  preparing  complex  metal-ammonia  compounds,  i.e. 
by  the  combination  of  metal  salt  and  ammonia  in  aqueous  solution,  has  oeen 
up  to  the  present  time  the  one  most  used.  It  should  be  noted,  however,  tiiat 
the  principal  form  in  which  the  complex  salt  exists  in  the  solution  is  not 
necessarily  that  form  in  which  it  crystallizes  out.  The  determination  of  the 
stoichiometric  composition  of  dissolved  metal-ammonia  salts  in  the  presence 
of  an  excess  of  ammonia  is  possible  by  a  combination  of  different  physico- 
chemical  methods  (solubility  measurements,  and  determinations  of  ion  con- 
centrations by  measuring  the  electromotive  force).  Thus  it  has  been  found 
that  the  composition  of  the  silver-ammonia  ion  which  preponderates  in 
aqueous  solution  is  [Ag(NH3)Ut,  whereas  the  salt  separating  from  such  a  solu- 
tion has  the  formula  2  AgCl.  3^H3. 

The  question,  with  which  of  the  components  of  the  salt  is  the  ammonia 
combined,  can  be  answered  by  the  physico-chemical  behavior  of  the  com- 
pound (conductivity  and  transference  number)  and  also  in  a  purely  chemical 
way.  A  solution  of  a  cupric  salt  to  which  an  excess  of  ammonia  is  added 
possesses  a  deep-blue  color  instead  of  the  usual  light-blue  of  copper  ions,  and 
the  addition  of  sodium  hydroxide  solution  no  longer  causes  a  precipitation  of 
cupric  hydroxide.  From  this,  and  from  the  fact  that  electrolysis  of  the  solu- 
tion causes  copper  to  migrate  as  before  towards  the  negative  pole,  the  con- 
clusion is  drawn  that  the  copper  ions  have  disappeared  and  been  replaced  by 
new  cations  containing  copper,  and  that  these  can  have  been  formed  only  by  the 
adding  of  ammonia  to  the  copper.  The  investigation  of  other  metal-ammonia 
salts  has  led  to  corresponding  results :  ammonia  is  usually  found  in  the  cation.1 

Whether  other  components  than  ammonia  take  part  in  the  formation  of 
the  complex  ions  is  ascertained,  first  by  conductivity  measurements,  from 
which  the  number  of  ions  contained  in  the  solution  is  determined,  and 
second  by  chemical  reactions  from  which  the  composition  of  the  ions  can  be 
deduced.  If  the  molecular  conductivity,  a,  of  1/512  normal  or  1/500  normal 
solutions  at  25°  is  compared  as  follows: 2 

Co(NH3)6X3  Co(NH3)5X3  Co(NH3)4X3  Co(NH3)3X3 

a        402  245  117  7 

and  at  the  same  time  the  conductivity  of  the  following  types  of  simple  salts 
under  the  same  conditions: 

Na3PO4  MgCl3  NaCl 

a          370  249  125 


1  For  exceptions  to  this  general  rule  see  No.  115  and  p.  169  (top). 

2  These    values     are     taken     from     measurements     on     [Co(NH3)6]Br3, 
[Co(NH3)5Br]Br2,[Co(NH3)4Cl2]Br,and[Co(NH3)3(NO2)3]WernerandMiolati; 
Z.  phys.  Chem.  12,  35  (1893);  14,  506  (1894);  21,  225  (1896). 


METAL-AMMONIA    COMPOUNDS.  165 

13  taken  into  consideration,  it  follows  that  the  first  of  the  above  complex 
Baits  must  be  a  quaternary  electrolyte,  the  second  a  ternary  one,  the  third  a 
binary  one,  while  the  fourth  is  not  an  electrolyte  at  all;  this  is  expressed  by 
the  following  formulas: 

[Co(NH3)6]X3        [Co(NH3)sX]X2    [CoCNH-^XJX         [Co(NH3)3XJ 

A  confirmation  of  this  view  is  furnished  by  the  chemical  reactions  of 
the  compounds.  If,  for  example,  X  is  a  halogen,  then  in  a  solution  of  the 
compound  [Co(NH3)5Br]Br2  only  two-thirds  of  the  bromine  is  precipitated 
by  the  addition  of  silver  nitrate.  In  a  similar  way  it  has  been  found,  in 
numerous  cases,  that  the  deductions  from  conductivity  measurements  are  in 
accord  with  the  chemical  behavior  of  substances,  and  thus  it  is  shown  that 
in  addition  to  ammonia,  also  certain  atoms  and  radicals  which  usually  act  as 
anions  (e.g.  the  halogens,  NO2,  and  CO3)  may  participate  in  the  formation  of 
complex  cations. 

In  one  series  of  metal-ammonia  salts  the  compounds  as  they  exist  in  the 
solid  state  contain  water  which  cannot  be  removed  without  destroying  the 
compounds.  This  water  remains  as  a  constituent  of  the  salts  even  when 
their  negative  components  are  replaced  by  others,  and  it  must  therefore 
belong  to  the  cation  (cf.  No.  134) 

[Co(NH3)5H2O]Cl3  +  3AgNO3  =  [Co(NH3)5H2O](NO3)3  +  3  AgCl. 

The  water  apparently  plays  the  same  part  in  these  compounds  as  ammonia 
in  the  first-mentioned  series,  as  is  evident  from  the  fact  that  the  conductivity 
is  not  materially  changed  when  molecules  of  ammonia  in  one  of  the  above 
compounds  are  replaced  successively  by  water  molecules: 

[Co(NH3)6]Br3         [Co(NH3)5H20]Br3  [Co(NH3)4(H2O)2]Br8 

a  402  390  380 

Besides  ammonia  a  number  of  other  substances  of  similar  nature,  such  as 
ethylene  diamine  and  pyridine,  often  take  part  in  the  formation  of  complex 
cations. 

A  comparison  of  the  composition  of  a  large  number  of  complex  com- 
pounds which  have  been  carefully  studied  from  different  points  of  view  has 
shown  that  here,  as  in  the  case  of  the  complex  anions,  the  number  of  atoms 
or  atomic  groups  which  are  attached  to  a  single  metal  atom  is  in  a  majority 
of  cases  either  four  or  six.  Werner1  accounts  for  this  by  assuming  here  also 
that  the  constituents  forming  the  complex  are  situated  around  the  central 
metal  atom  either  in  the  four  corners  of  a  square  or  in  the  six  corners  of  an 
octahedron.  The  ever-recurring  co-ordination  numbers  4  and  6  would  thus 
be  a  measure  of  the  available  space  around  the  central  atom,  but  they  are  quite 
independent  of  the  valence  of  this  atom.  This  space,  which  Werner  desig- 
nates as  the  "inner  sphere"  is  characterized  by  the  fact  that  no  components 
can  be  dislodged  from  it  by  electrolytic  dissociation  and  that  the  whole  is 
able  to  take  part  in  chemical  reactions  exactly  in  the  same  way  as  an  indi- 

1  Neuere  Anschauungen  auf  dem  Gebiete  der  anorganischen  Chemie,  2nd 
Ed.  Braunschweig  1909. 


166  METAL-AMMONIA    COMPOUNDS. 

vidual  metal  atom.  The  valence  of  the  complex  is  equal  to  the  difference 
between  that  of  the  central  atom  and  the  valence  of  the  attached  acid  radi- 
cals. If  the  valences  of  the  acid  residues  are  sufficient  to  satisfy  that  of  the 
central  atom,  then  the  total  valence  of  the  complex  is  zero,  i.e.  the  sub- 
stance is  a  non-electrolyte,  e.g.  trinitrito-triammine-cobalt.  (No.  139.) 

Nomenclature.  —  Before  the  constitution  of  the  metal-ammonia  com- 
pounds was  known  certain  series  were  designated  according  to  their  external 
characteristics  (e.g.  color,  in  roseo  salts  and  purpureo  salts)  while  others 
were  named  after  the  discoverer  (e.g.  Magnus'  salt).  Werner,  after  formu- 
lating his  theory,  proposed  a  rational  system  of  nomenclature  which  is  now 
quite  generally  adopted.  According  to  this  system  the  names  of  the  con- 
stituents belonging  to  the  complex  cation  are  placed  before  the  name  of  the 
metal  and  are  arranged  in  the  following  order:  first  the  acid  radicals;  then 
any  groups  which  behave  in  the  same  manner  as  ammonia;  and  lastly,  just 
preceding  the  metal,  the  ammonia  itself.  The  acid  radicals  are  given  the 
suffix  "o"  while  ammonia  is  designated  by  the  term  " ammine,"  and  water 
contained  in  the  complex  is  designated  as  aquo.  The  number  of  each  kind  of 
constituent  is  indicated  by  the  appropriate  prefix,  di,  tri,  etc.  The  com- 
plete name  of  the  cation  is  commonly  printed  as  a  single  word.  The  anions  are 
designated  in  the  usual  way  and  follow  the  name  of  the  cation.  Further 
particulars  will  be  evident  from  a  study  of  the  examples  in  the  accompany- 
ing tables,  pp.  167  and  168. 

Isomerism.  It  is  apparent  from  the  tables  on  the  following  two  pages 
that  in  certain  cases  two  compounds  exist  which  have  the  same  empirical 
composition,  and  the  same  molecular  weight,  but  still  have  different  prop- 
erties (isomerism).  The  existence  of  such  isomers  leads  here,  as  in  all 
similar  cases,  to  the  assumption  that  there  is  a  definite  spatial  arrangement 
of  the  various  parts  of  the  molecule  and  gives  justification  for  the  arbi- 
trary assumption  that  the  constituents  of  the  "inner  sphere"  which  sur- 
rounds the  central  metal  atom  are  actually  situated  at  the  corners  of  a  square 
or  of  an  octahedron.  Two  different  pairs  of  substituents  are  manifestly 
capable  of  two  different  arrangements  at  the  corners  of  a  square: 

NH,V       /Cl  NH3V       /Cl 

;pt(  pt 

NH3X       XC1  Gl<       XNH3 

1.2  or  cis-position  1.3  or  trans-position. 

If,  however,  three  substituents  in  a  square  are  the  same  and  the  fourth 
different,  such  an  isomerism  is  impossible.  The  actual  lack  of  isomers  in  the 
latter  case  and  their  existence  in  the  former,  strongly  supports  the  above 
assumption. 

Similarly,  the  isomerism  of  the  two  dinitritotetramminecobaltic  series  is 
explained  by  the  use  of  the  corresponding  octahedral  formulas  (cf.  No.  138). 

As  has  already  been  stated,  the  external  valence  of  the  complex  is  equal  to 
the  difference  between  the  valence  of  the  metal  and  the  total  valence  of  the 
acid  radicals  contained  in  the  complex.  This  external  valence  may  become 
zero,  or  even  negative  when  the  valence  of  the  acid  part  of  the  complex 


METAI^AMMONIA  COMPOUNDS. 


167 


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SILVER-AMMONIA  SULPHATE.  169 

preponderates.     In  the  following  series  of  seven  cobalt  compounds,  the  valence 
of  the  complex  changes  progressively  from  three  positive  to  three  negative : 
[Co(NH3)  6]  Cl,     [Co(NH3)  5N02]  C12      [Co(NH3)4(NO2)2]  Cl     [Co(NH3)  3(NO2)  J 
K  [Co(NH3)2(N02)4]  (No.  115)  K2  [Co(NH3)(NO2)5] 

K3[Co(N02)6]  (No.  114). 

The  binding  power  for  potassium,  or  in  other  words  the  affinity  for  nega- 
tive electrons,  is  occasioned  by  the  valences  of  the  nitrite  radicals.  The  bind- 
ing of  these  nitrite  groups  within  the  complex  takes  place  in  part  by  means  of 
the  secondary  valences  (cf.  pp.  141  and  155). 


NO2-Co 


..,N02-1-K 
-NO2-  -K 
XNO2-J-K 


so  that  of  the  six  nitrite  groups  attached  to  the  cobalt,  three  are  held  by  the 
principal  valences  and  three  by  the  secondary  valences  of  the  metal.  Inas- 
much as  the  secondary  valences  are  weaker  than  the  principal  ones,  this  view 
corresponds  well  with  the  formation  of  potassium  cobaltinitrite  from  potas- 
sium nitrite  and  cobaltic  nitrite,  and  with  its  manner  of  dissociating.  For 
the  further  development  of  this  theory,  which  leads  to  the  conception  of 
"indirect  combination,"  the  work  of  Werner  already  cited  should  be  con- 
sulted. 

125.   Silver-ammonia  Sulphate  [Ag(NH3)2]2  SO4. 

Diamminesilver  sulphate  can  be  prepared  in  the  solid  form  although  this 
is  not  possible  with  the  corresponding  chloride  (cf.  p.  164). 

Treat  10  g.  of  silver  nitrate  with  31  c.c.  of  2-normal  sulphuric 
acid  and  heat  the  mixture  in  a  small  evaporating  dish  on  an  air 
bath,  until  no  more  acid  vapors  are  given  off.  While  still  hot 
dissolve  the  perfectly  dry  residue  in  the  least  amount  possible  of 
concentrated  ammonia  (Hood).  Filter  the  solution  and  allow  it 
to  crystallize.  Recrystallize  the  crude  product  from  water  con- 
taining a  little  ammonia.  Yield,  about  5  grams  of  colorless, 
columnar  crystals. 

Analyze  the  product  by  determining  the  silver  as  AlgCl  and  the 
ammonia  by  distilling  with  caustic<  soda  and  titrating  the  dis- 
tillate. 

126.   Tetramminecupric  Sulphate,  [Cu(NH3)4]  SO4.H2O  and  Ammo- 
nium Cupric  Sulphate,  CuSO4.(NH4)2  SO4.6H2O. 

The  first  of  these  salts  may  be  regarded  as  ordinary  blue  vitriol  in  which 
four  molecules  of  water  have  been  replaced  by  ammonia.  The  fifth  mole- 
cule of  water  evidently  has  another  function  from  that  of  the  other  four, 
cf.  p.  189  (top). 


170  METAL-AMMONIA  COMPOUNDS. 

Dissolve  50  g.  of  powdered  copper  vitriol  in  a  mixture  of  75  c.c. 
concentrated  ammonia  and  50  c.c.  of  water  and  precipitate  the 
deep-blue  solution  (first  filtering  it  through  asbestos  if  necessary) 
by  the  gradual  addition  of  75  c.c.  alcohol.  After  standing  for 
some  hours  in  the  cold,  filter  off  the  dark-blue  crystals  with  suc- 
tion, and  wash  them  first  with  a  mixture  of  equal  volumes  of 
alcohol  and  concentrated  ammonia  and  finally  with  alcohol  and 
ether.  Dry  the  sajt  at  the  laboratory  temperature.  Yield, 
almost  quantitative. 

Ammonium  Cupric  Sulphate  crystallizes  on  cooling  from  a  hot 
solution  of  25  g.  copper  vitriol  and  13.25  g.  ammonium  sulphate 
in  40  c.c.  of  water.  Recover  the  last  portions  of  the  salt  by  con- 
centrating the  mother-liquor.  The  yield  is  almost  quantitative, 
and  consists  of  large,  light-blue  crystals.  The  aqueous  solution 
of  the  salt  gives  a  precipitate  of  cupric  hydroxide  on  being 
treated  with  sodium  hydroxide.  The  salt  belongs  to  the  large 
class  of  monoclinic  double  salts  which  crystallize  with  six  mole- 
cules of  water. 

127.   Tetramminecupric  Chloride,  [Cu(NH3)4]Cl2.H2O. 

In  order  to  prepare  this  salt,  which  is  readily  soluble  in  water, 
conduct  ammonia  gas  through  a  wide  delivery  tube  into  a  warm 
filtered  solution  of  17  g.  crystallized  cupric  chloride  in  15  c.c.  of 
water.  The  ammonia  gas  is  obtained  by  heating  100  c.c.  of  con- 
centrated ammonia  solution  gently  and  passing  the  gas  through 
an  empty  wash-bottle  to  free  it  to  some  extent  from  water  vapor. 
The  mixture  becomes  heated  to  boiling  as  a  result  of  the  reaction, 
and  it  should  therefore  be  cooled  somewhat  to  prevent  too  much 
evaporation.  When  a  clear,  deep-blue  solution  is  obtained  add 
8  c.c.  of  alcohol  and  again  saturate  the  liquid  with  ammonia,  this 
time  cooling  it  with  ice.  Drain  the  deposited  salt  on  a  hardened 
filter,  wash  it  with  alcohol  to  which  a  little  concentrated  ammonia 
has  been  added,  then  with  pure  alcohol  and  finally  with  ether, 
and  suck  it  as  dry  as  possible.  Yield,  15  to  18  grams.  The 
preparation  must  be  placed  immediately  in  a  stoppered  bottle, 
as  it  loses  ammonia  on  .standing  in  the  open  air.  The  addition 
of  the  alcoholic  washings  to  the  mother-liquor  causes  the  pre- 
cipitation of  a  few  grams  more  of  finely-divided  and  consequently 
lighter-colored  product. 


CARBONATOTETRAMMINECOBALTIC   NITRATE.  171 


128.   Hexamminenickelous  Bromide,  [Ni(NH3)6]Br2. 

Hexamminenickelous  bromide  can  be  used  in  making  cobalt-free  prepa- 
rations of  nickel  such  as  are  required  for  atomic  weight  determinations. 
The  corresponding  cobalt  salt  is  far  more  soluble.  The  raw  material  used 
and  the  final  product  obtained  should  be  tested  qualitatively  for  cobalt. 

Pour  a  solution  of  141  g.  crystallized  nickel  sulphate  in  4  liters 
of  water  into  a  large  flask  and  add  a  solution  of  41  g.  sodium 
hydroxide  in  200  c.c.  of  water.  Wash  the  resulting,  voluminous, 
light-green  precipitate  by  decantation;  let  it  settle  as  much  as 
possible,  then  siphon  off  the  supernatant  liquid  and  fill  the  flask 
again  with  distilled  water;  repeat  this  process  from  time  to  time 
for  from  three  to  five  days.  Collect  the  washed  precipitate  upon 
a  large  plaited  filter  and  dissolve  it  in  an  aqueous  solution  of  hydro- 
bromic  acid  containing  about  82  grams  of  the  acid.1  Evaporate 
the  filtered  solution  to  dryness  on  the  water-bath,  take  up  the 
residue  in  as  little  water  as  possible,  cool  the  solution  to  0°  with 
a  mixture  of  ice  and  salt  and  then  treat  it  with  concentrated 
ammonia;  an  abundant  precipitate  of  violet,  micro-crystalline 
flakes  is  obtained  and  the  liquid  becomes  colorless.  After  stand- 
ing for  a  short  time  at  0°,  collect  the  crystalline  mass  on  a  suction- 
filter,  wash  it  carefully  with  ice-cold  ammonia  (to  remove  any 
cobalt  present),  and  dry  it  in  a  desiccator  over  lime  which  is 
mixed  with  a  little  solid  ammonium  chloride.  The  yield  is  almost 
theoretical.  If  desired  the  preparation  can  be  carried  out  on  a 
smaller  scale. 

129.    Carbonatotetramminecobaltic  Nitrate, 
[Co(NH3)4C03]  N03.JH20. 

When  cobalt  salts  are  treated  out  of  contact  with  the  air  with  an  excess  of 
ammonia,  hexamminecobaltous  salts  are  formed.  If  such  a  solution  is 
mixed  with  a  large  amount  of  ammonium  carbonate,  oxidized  by  the  air,  and 
then  evaporated  to  a  small  volume  with  the  addition  of  more  ammonium 
carbonate,  crystals  of  Carbonatotetramminecobaltic  salt  are  produced. 

Dissolve  20  g.  of  cobalt  carbonate  in  as  little  concentrated 
nitric  acid  as  possible  and  dilute  the  solution  to  100  c.c.  (or  50  g. 
of  crystallized  cobaltous  nitrate  may  be  dissolved  in  an  equal 
amount  of  water). 

Cf.  No.  35. 


172  .  METAL-AMMONIA    COMPOUNDS. 

Dissolve  100  g.  of  ammonium  carbonate  in  500  c.c.  of  water  in 
a  1-liter  flask,  add  250  c.c.  of  concentrated  ammonia,  and  into 
this  solution  pour  the  above  prepared  solution  of  cobalt  nitrate. 
Draw  a  stream  of  air  through  the  deep-violet-colored  liquid  in  the 
manner  described  under  No.  107.  At  the  end  of  three  hours, 
pour  the  liquid,  which  is  now  of  a  blood-red  color,  into  a  porce- 
lain evaporating  dish,  and  concentrate  it  to  about  300  c.c.,  adding 
5  g.  of  powdered  ammonium  carbonate  every  fifteen  minutes 
(about  25  or  30  grams  in  all).  Filter  the  solution  immediately 
from  any  small  amount  of  sediment  and  boil  it  down  to  200  c.c., 
still  continuing  to  add  ammonium  carbonate  in  small  quantities 
(about  10  g.  in  all  during  this  final  evaporation).  Allow  the 
liquid  to  cool  and  collect  the  large,  purple  crystal-plates  which 
separate  out.  Wash  the  crystals  with  a  little  water,  then  with 
dilute  alcohol  and  finally  with  pure  alcohol./  Recover  a  crop  of 
impure  crystals 'by  concentrating  the  mother-liquor;  extract  this 
product  at  the  laboratory  temperature  with  15  times  its  weight 
of  water,  and  precipitate  the  salt  from  the  filtered  solution  by 
gradually  adding  two  to  three  volumes  of  alcohol.  Collect 
this  precipitate  and  wash  it  as  directed  above.  Total  yield,  22  to 
25  grams. 

130.    Chloropentamminecobaltic    Chloride,  [Co(NH3)5Cl]Cl2,  from 
the  preceding. 

The  preparation  of  Chloropentamminecobaltic  chloride  from  the  carbon- 
atotetrammine  salt  is  an  example  of  the  replacement  of  certain  constituents 
in  the  complex  by  means  of  others.  On  acidifying  an  aqueous  solution  of 
the  carbonatotetrammine  salt  with  hydrochloric  acid,  there  is  formed,  for  the 
most  part,  chloroaquotetramminecobaltic  chloride.  This  is  changed  on  being 
heated  in  an  ammoniacal  solution,  into  aquopentamminecobaltic  chloride,  and 
the  latter  on  being  acidified  with  hydrochloric  acid  is  converted  into  Chloro- 
pentamminecobaltic chloride.  This  is  one  of  the  best  and  longest-known  mem- 
bers of  the  entire  series,  and  has,  on  account  of  its  color  and  chlorine  content, 
been  called  chloropurpureocobaltic  chloride. 

Add  concentrated  hydrochloric  acid  (about  4.5  c.c.)  to  a  solution 
of  3  g.  carbonatotetramminecobaltic  nitrate  in  40  c.c.  of  water 
until  all  the  carbon  dioxide  has  been  expelled.  Then  make  the 
solution  slightly  ammoniacal;  thereupon  add  an  excess  of  5  c.c. 
concentrated  ammonia  and  heat  the  solution  for  three-quarters 
of  an  hour  on  the  water-bath.  After  cooling  add  50  c.c.  of  con- 


CHLOROPENTAMMINECOBALTIC  CHLORIDE  173 

centrated  hydrochloric  acid  and  heat  the  mixture  once  more  for 
an  hour  on  the  water-bath.  -Drain  the  violet-red  salt  which 
separates  and  wash  it  with  alcohol.  Yield,  1.5  ,to  2.5  grams. 

131.    Chloropentamminecobaltic  Chloride  ^rom  Cobalt  Carbonate. 

By  oxidizing  a  strongly  ammoniacal  solution  of  hexamminecobaltous 
chloride  containing  a  large  amount  of  ammonium  carbonate,  carbonato- 
tetramminecobaltic  chloride  (cf.  No.  129),  aquopentammine  cobaltic  chloride, 
[Co(NH3)6H2O]Cl3,  and  oxycobaltammine  chloride  [Co2O2(NH3)10]Cl4  are 
formed.  By  adding  ammonium  chloride  in  considerable  quantity  to  the  solu- 
tion and  evaporating,  the  last  mentioned  compound  is  converted  into  chloro- 
pentamminecobaltic  chloride,  or  into  'aquopentamminecobaltic  chloride. 
These  substances,  like  carbonatotetramminecobaltic  chloride,  can  be  con- 
verted into  Chloropentamminecobaltic  chloride  (cf.  preceding  preparation). 

Dissolve  20  g.  of  cobalt  carbonate  in  as  little  hydrochloric  acid 
as  possible  and  treat  the  filtered  solution  in  the  cold  with  250  c.c. 
of  10%  ammonia  and  50  g.  of  ammonium  carbonate  dissolved  in 
250  c.c.  of  water.  Oxidize  the  mixture  by  passing  through  it  a 
rapid  current  of  air  for  three  hours  (cf.  No.  107).  Then  add  150  g. 
of  ammonium  chloride  and  evaporate  the  solution  on  the  water- 
bath  until  it  becomes  of  pasty  consistency.  Acidify  with  hydro- 
chloric acid,  while  stirring,  until  no  more  carbon  dioxide  is  evolved 
and  then  make  ammoniacal  once  more,  adding  10  c.c.  of  con- 
centrated ammonia  in  excess.  After  diluting  to  400-500  c.c. 
heat  the  mixture  again  on  the  water-bath  for  an  hour.  Add  300  c.c. 
of  concentrated  hydrochloric  acid  ahd  leave  the  solution  on  the 
water-bath  until,  at  the  end  of  from  half  to  three-quarters  of  an 
hour,  a  separation  of  Chloropentamminecobaltic  chloride  takes 
place.  After  cooling,  filter  off  the  crystals  and  wash  them  with 
dilute  hydrochloric  acid. 

For  purification,  dissolve  the  crude  product  in  300  c.c.  of  two 
per  cent  ammonia  solution,  whereby  aquopentamminecobaltic 
chloride  is  formed  (cf.  No.  134),  extract  the  residue  twice  with 
50  c.c.  of  the  same  ammonia  solution,  and  precipitate  the  filtrate 
by  adding  300  c.c.  of  concentrated  hydrochloric  acid  and  heat- 
ing the  mixture  for  three-quarters  of  an  hour  on  the  water-bath. 
After  it  has  become  perfectly  cold,  filter  off  the  salt,  and  wash  it 
with  dilute  hydrochloric  acid  and  alcohol.  Yield,  30  to  34  grams, 
which  is  nearly  the  theoretical  quantity. 

Since  nickel  does  not  form  any  similar  salt,  this  gives  a  means  of  preparing 
a  nickel-free  product  from  impure  cobalt  preparations. 


174  METAL-AMMONIA    COMPOUNDS. 

Reactions.  A  cold,  concentrated,  aqueous  solution  of  the  above 
salt  when  treated  with  hydrochloric  acid  yields  short,  red  crystals 
after  standing  some  time.  Mercuric  chloride  causes  the  imme- 
diate precipitation  of  long,  rose-red  needles  of  the  composition 
[Co(NH3)5Cl]Cl2.3  HgCl2.  Potassium  chromate  gives  brick-brown 
crystals.  Potassium  pyrochromate  causes  the  slow  formation  of 
clusters  of  fine,  orange  needles,  or  of  large  flat  prisms;  ammonium 
oxalate  slowly  yields  well-formed,  red  prisms  which  can  be  seen 
under  the  microscope. 

132.   Sulphate  and  Nitrate  of  the  Chloropentammine  Series. 

Acid  Chloropentamminecobaltic  Sulphate,  [Co(NH3)5CJ]2(HS04)2S04. 
Triturate  5  g.  of  Chloropentamminecobaltic  chloride  with  12  g. 
of  concentrated  sulphuric  acid  in  a  mortar,  whereby  hydrogen 
chloride  is  evolved.  Dissolve  the  mass  in  40  c.c.  of  water  at 
70°  and  filter.  On  cooling  the  solution,  long,  thin  fuchsine-red 
prisms  are  deposited.  The  crystals  become  larger  if  they  are 
allowed  to  stand  for  several  days  in  contact  with  the  mother- 
liquor  because  smaller  crystal  grains  are  slightly  more  soluble 
than  larger  ones  and  therefore  dissolve  slowly,  while  the  larger 
crystals  grow  still  larger.  Wash  the  crystals  with  absolute 
alcohol  and  dry  them  in  the  hot  closet.  Yield,  3  grams. 

When  the  alcoholic  washings  are  added  to  the  mother-liquor 
a  fine,  dull-red,  crystalline  meal  of  the  same  salt  is  precipitated. 
Dissolve  it  in  a  little  warm  water,  cool  the  solution  and  precipitate 
the  nitrate  by  adding  concentrated  nitric  acid  (Mass-action). 
Recrystallize  in  order  to  obtain  a  purer  product.  Small  dull-red 
crystals. 

Reactions.  The  solution  of  the  pure  sulphate  gives  no  precipitate 
when  treated  with  silver  nitrate  at  the  room  temperature;  on 
heating,  however,  some  silver  chloride  is  thrown  down,  but  even 
on  boiling,  the  reaction  is  incomplete.  On  the  other  hand,  the 
sulphate  radical,  not  being  contained  in  the  complex,  gives  at 
once  a  quantitative  precipitation  with  barium  chloride.  •  The 
nitrate  shows  a  similar  behavior  towards  silver  ions;  the  solution 
becomes  turbid  only  on  boiling. 

In  the  analysis  of  the  salt,  chlorine  and  sulphate  are  determined 
after  the  substance  has  been  decomposed  by  boiling  with  caustic 
soda;  cobalt  is  weighed  as  sulphate  after  igniting  the  salt  and 


HEXAMMINECOBALTIC  SALTS.  175 

heating  the  residue  with  concentrated  sulphuric  acid;  ammonia 
is  estimated  in  the  usual  manner  by  distilling  with  caustic  soda 
and  titrating  the  distillate. 

133.   Hexamminecobaltic  Salts,  [Co(NH3)6]  X3  (Luteocobalt  Salts). 

1.  Luteocobalt  Chloride.  Place  10  g.  of  chloropentammine- 
cobaltic  chloride,  8  g.  of  ammonium  chloride,  and  100  c.c.  of  20% 
ammonia  solution  in  a  soda-water  bottle  (pressure-flask),  and 
stopper  it  tightly.  Wrap  the  bottle  firmly  in  a  towel,  fasten  it  to 
a  wooden  handle  and  heat  it  for  six  hours  in  a  pail  of  boiling  water, 
shaking  thoroughly"  once  every  two  hours.  Since  a  strong  pres- 
sure prevails  in  the  bottle,  this  part  of  the  process  should  be 
carried  out  outside  of  the  general  laboratory,  and  before  each 
shaking  the  mixture  should  be  removed  from  the  bath  and  allowed 
to  cool  somewhat.  At  the  end  of  the  heating  the  chloropent- 
ammine  salt  must  have  disappeared  almost  entirely.  When 
cold  open  the  bottle,  pour  the  contents  into  an  evaporating  dish, 
and  let  it  stand  24  hours  in  the  open  air,  or  under  the  hood,  to 
allow  the  ammonia  to  volatilize.  Dilute  with  300  to  400  c.c.  of 
water,  add  50  c.c.  of  concentrated  hydrochloric  acid  and  heat  the 
mixture  in  a  flask  for  an  hour  on  the  water-bath.  Add  250  c.c. 
more  of  concentrated  hydrochloric  acid  and  cool  rapidly,  while 
shaking,  under  the  wrater-tap.  Drain  the  yellow  precipitate  on 
a  hardened  filter  paper  and  wash  it  with  20%  hydrochloric 
acid. 

To  purify  the  first  product,  dissolve  it  in  as  little  cold  water  as 
possible  (about  5  g.  of  the  chloride  dissolves  in  100  c.c.  of  water), 
filter  the  solution  from  any  residual  chloropentamminecobaltic 
chloride,  and  precipitate  the  yellow  filtrate  (keeping  the  liquid 
cold)  by  the  gradual  addition  of  half  its  volume  of  concentrated 
hydrochloric  acid.  Yield,  8  to  10  grams. 

Reactions.  A  cold,  saturated,  aqueous  solution  of  the  luteo- 
chloride  when  treated  with  ammonium  oxalate  yields  the  very 
difficultly  soluble  oxalate  which  consists  of  small,  light-brownish- 
yellow,  irregular  crystals.  With  mercuric  chloride  a  voluminous, 
light-pink  double  salt,  [Co(NH3)6]Cl3.3HgCl2.H2O,  is  at  once 
thrown  down.  Potassium  chromate  immediately  precipitates 
brownish-yellow  clusters  of  needles.  Potassium  pyrochromate 


176  METAL-AMMONIA  COMPOUNDS. 

causes  the  immediate  separation  of  a  precipitate  which  under  the 
nicroscope  is  seen  to  be  crystalline. 

2.  Luteocobalt  Nitrate  by  the  Iodine  Method.  Dissolve  24  g. 
of  cobalt  carbonate  by  warming  it  with  a  just  sufficient  amount  of 
dilute  nitric  acid,  filter,  and  dilute  the  solution  to  100  c.c.  Add 
200  c.c.  of  concentrated  ammonia,  heat  the  solution  to  boiling, 
and  oxidize  the  salt  by  adding  25.4  g.  of  iodine,  which  must  be 
introduced  slowly  at  first.  A  vigorous  reaction  takes  place  and  a 
pale  yellowish-brown  precipitate  of  luteo  salt  is  formed.  All  the 
iodine  should  be  added  in  the  course  of  half  an  hour.  Allow  the 
liquid  to  cool  and  after  it  has  stood  about  two  hours  filter  off 
the  precipitate  and  wash  it  with  water  containing  ammonia.  Then 
boil  the  salt  with  200  c.c.  of  approximately  56%  nitric  acid,  where- 
by iodine  is  set  free  which  can  be  recovered  to  some  extent  by 
means  of  two  funnels,  one  placed  in  the  flask  and  the  other  inverted 
over  it  to  form  a  double  cone.  When  all  the  iodine  has  been 
expelled,  filter  off  the  precipitate,  drain  it  with  suction,  wash  it 
with  water  containing  nitric  acid  and  finally  with  alcohol,  and  dry 
it  in  the  hot  closet.  Yield,  about  22  grams. 

134.   Aquopentamminecobaltic  Salts,  [Co(NH3)5H2O]  C13  (Roseo- 

cobalt  Salts). 

Chloropentamminecobaltic  chloride  dissolves  in  ammonia,  forming  aquo- 
pentamminecobaltic  chloride.  If  this  solution  is  precipitated  warm  with 
hydrochloric  acid,  the  chloropentammine  salt  is  formed  again,  but  if  the 
hydrochloric  acid  is  added  slowly  and  the  solution  kept  very  cold,  aquopent- 
amminecobaltic  chloride  separates.  Less  care  need  be  taken  in  the  prepa- 
ration of  the  difficultly  soluble  aquopentamminecobaltic  oxalate.  Aquopent- 
ammine  salts  are  also  formed  by  the  oxidation  of  hexamminecobaltous  salts 
by  potassium  permanganate. 

1.  Dissolve  10  g.  of  Chloropentamminecobaltic  chloride  in  a 
liter  flask  by  shaking  it  with  300  c.c.  of  5%  ammonia  and  heating 
upon  the  water-bath.  Filter,  if  necessary,  and  then  cool  the  solu- 
tion to  about  0°,  first  by  holding  the  flask  under  running  water 
and  then  by  surrounding  it  with  ice.  While  rotating  the  flask  and 
keeping  its  contents  cold,  add  strong  hydrochloric  acid  from  a 
dropping-funnel,  a  few  drops  at  a  time,  until  the  solution  reacts 
acid.  Filter  and  drain  the  bright-red,  crystalline  precipitate; 


AQUOPENTAMMINECOBALTIC  SALTS.  177 

wash  it  first  with  a  little  50%  alcohol  then  with  pure  alcohol 
and  dry  it  in  a  fairly  warm  place.     Yield,  10  grams. 

2.  To  a  cold  solution  of  20  g.  crystallized  cobaltous  chloride 
in  360  c.c.  of  water,  contained  in  a  1500  c.c.  flask,  add  110  c.c.  of 
concentrated   ammonia   and   then    10   g.   of  potassium  perman- 
ganate  dissolved  in   400   c.c.   of  water.     Shake  the   mixture   a 
number  of  times,  and,  after  it  has  stood  for  24  hours,  filter  off  the 
slime   of   hydrated   manganese   dioxide.     Neutralize   the   filtrate 
with  dilute  hydrochloric  acid,  and  while  keeping  it  cold  by  surround- 
ing the  vessel  with  ice,  precipitate  the  product  by  the  gradual 
addition  of  a  mixture  of  three  volumes  of  concentrated  hydro- 
chloric acid  and  one  volume  of  alcohol.     Wash  the  precipitate 
with  alcohol.     Yield,  12  to  15  grams. 

3.  Purification   of  Aquopentamminecobaltic    Chloride.    Dissolve 
the  crude  product,  prepared  according  to  either  1  or  2,  in  cold, 
2%  ammonia,  using  75  c.c.  for  each  10  g.  of  the  salt;  filter  off  the 
slight  residue  of  luteo  salt,  and,  while  keeping  the  solution  cold 
with  ice,  precipitate  the  roseo  salt  by  the  gradual  addition  of  con- 
centrated  hydrochloric   acid.     Drain   the   precipitate   and   wash 
it  with  a  mixture  of  equal  parts  concentrated  hydrochloric  acid 
and  water,  then  with  alcohol,  and  dry  it  in  a  warm  place. 

4.  Aquopentamminecobaltic   Oxalate.    Dissolve  10  g.  of  chloro- 
pentamminecobaltic  chloride  in  a  flask  with  75  c.c.  of  water  and 
50  c.c.  of  10%  ammonia,  heating  on  the  water-bath.     After  cool- 
ing the  deep-red  solution  to  the  room  temperature,  filter  it  and 
treat  with  a  solution  of  oxalic  acid  until  a  precipitate  begins  to 
form,  then  continue  to  add  the  reagent  very  carefully  until  the 
solution  is  just  acid.     Filter,  and  \vash  the  salt  with  water. 

To  purify  the  product,  dissolve  it  as  directed  urider  3,  and  re- 
precipitate  the  salt  by  the  careful  addition  of  oxalic  acid  solution. 
Drain  the  precipitate,  wash  it  with  water,  then  with  alcohol,  and 
dry  it  in  a  warm  place.  Yield,  9  grams  of  pure1  substance. 

5.  Aquopentamminecobaltic   Chloride  from   Aquopentammineco- 
baltic Oxalate.      Cover  10  g.  of  aquopentamminecobaltic  oxalate 
at  room  temperature  with  30  c.c.  of  water  and  dissolve  the  salt  by 
the  addition  of  50  c.c.  of  normal  hydrochloric  acid.     Cool  the  solu- 


1  The  separation  of  luteo  salt  present  as  impurity  is  here  more  complete 
than  according  to  procedure  3,  because  its  oxalate  is  much  more  insoluble 
than  its  chloride. 


178  METAI^AMMONIA  COMPOUNDS. 

tion  with  ice  to  about  0°  and  precipitate  it  very  slowly  by  allow- 
ing 100  c.c.  of  concentrated  hydrochloric  acid  to  flow  upon  it 
drop  by  drop.  Drain  off  the  bright  red  precipitate  and  wash  it 
first  with  hydrochloric  acid  diluted  to  one-half  and  then  with 
alcohol.  Yield,  about  8  g.  of  perfectly  pure  aquopentammineco- 
baltic  chloride. 

135.   Dibromotetramminecobaltic    Bromide,  [Co(NH3)4Br2]  Br, 
(Dibrompraseo  Salt). 

Place  20  g.  of  carbonatotetramminecobaltic  nitrate  in  a  flask, 
shake  it  with  80  g.  of  concentrated  hydrobromic  acid  (cf.  No.  35) 
and  continue  to  shake  while  heating  slowly  over  a  free  flame. 
The  mass  evolves  a  large  amount  of  carbon  dioxide  and  changes 
at  first  to  reddish-brown,  then  pale-brown,  and  finally  to  a  dull- 
green.  When  the  color  ceases  to  change,  allow  the  mixture  to 
cool  to  room  temperature  and  add  50  c.c.  of  cold  water,  which 
serves  to  dissolve  a  little  admixed  bromoaquotetramminecobaltic 
bromide.  Drain  the  precipitate  on  a  hardened  filter- and  wash 
it  with  cold  water  until  the  filtrate  is  no  longer  reddish-violet  but 
comes  through  colorless.  Yield,  24  grams  of  a  very  fine,  light- 
yellowish-green,  difficultly-soluble  powder. 

Dissolve  a  portion  of  the  powder  in  warm,  dilute  ammonia, 
acidify  with  hydrobromic  acid,  and  heat  for  half  an  hour  on  the 
water-bath.  The  salt  is  thereby  transformed  into  bromopent- 
amminecobaltic  bromide,  [Co(NH3)5Br]Br2,  which  is  a  very 
finely  granular,  reddish- violet  precipitate. 

136.    1.2    Dinitritotetramminecobaltic    Salts,    [Co(NH3)4(NO2)2]  X 
(Flavocobalt  Salts.) 

Flavocobalt  Nitrate,  [Co(NHz)^(N02)2]N03.  Dissolve,  with- 
out heating,  10  g.  of  carbonatotetramminecobaltic  nitrate  in  a 
mixture  of  100  cc.  water  and  14  g.  of  40  %  nitric  acid  (sp.  gr. 
1.25),  whereby  a  blood-red  solution  of  diaquotetrarnminecobaltic 
nitrate  is  obtained.  Add  20  g.  of  crystallized  sodium  nitrite  to 
this  solution,  little  by  little,  and  place  the  flask  containing  the 
mixture  in  a  bath  of  boiling  water  for  7-8  minutes,  or  until  the 
color  of  the  solution  has  become  deep  brownish-yellow,  then 
immediately  cool  under  the  water  tap  and  add  130  cc.  more  of 
the  40  %  nitric  acid.  This  causes  foaming  and  a  strong  evolution 


DINITRITOTETRAMMINECOBALTIC  CHLORIDE.  179 

of  nitric  oxide.  The  next  morning,  drain  off  the  precipitate  of 
mixed  acid  and  neutral  flavonitrate  and  wash  it  first  with  a  little 
nitric  acid  and  then  with  alcohol.  Yield,  8-9  grams. 

To  purify  the  salt,  recrystallize  from  25  cc.  of  water  slightly 
acidulated  with  acetic  acid.  Collect  the  light-brown,  crystalline 
flakes,  or  prisms,  which  are  obtained,  and  wash  with  50  %  alco- 
hol, then  with  pure  alcohol,  and  dry  at  a  gentle  heat.  Yield,  6.5 
grams. 

Reactions.  A  cold,  saturated  solution  of  flavonitrate  (3  g.  in 
100  cc.  water)  gives  with  potassium  chromate  a  crystalline  pre- 
cipitate of  irregularly  indented  leaflets  which  are  crossed  and 
branched  at  right  angles  (microscope).  With  potassium  pyro- 
chromate  small  clusters  of  fine  needles  are  at  once  precipitated, 
which  are  inclined  towards  one  another  like  the  branches  of  a  fir 
tree. 

Flavocobalt  Chloride,  [Co(NH3)^(N02)2]CL  Dissolve  1  g.  of 
flavonitrate  in  30  cc.  of  water,  warming  gently;  add  2  g.  of 
ammonium  chloride  and  filter  if  necessary.  Then  gradually  add 
100  cc.  of  alcohol  to  the  mixture  and  after  standing  24  hours  filter 
off  the  small,  deep-yellow  crystal  leaflets.  Wash  the  product 
with  dilute  alcohol,  then  with  pure  alcohol,  and  dry  it  in  the  hot 
closet.  Yield,  0.9  gram. 

'l37.    1.6  Dinitritotetramminecobaltic  Chloride,  [Co(NH3)4(NO2)2]Cl 
(Croceocobalt  Chloride). 

To  a  cold,  filtered  solution  of  100  g.  ammonium  chloride  and 
135  g.  sodium  nitrite  in  750  cc.  water,  add  150  cc.  of  20  %  ammo- 
nia and  90  g.  of  cobalt6us~chloride  dissolved  in  250  cc.  of  water. 
Draw  air  through  this  solution  (cf.  No.  107)  for  four  hours,  whereby 
the  color,  which  is  at  first  a  weak,  brownish-green,  changes  to 
green  with  a  tinge  of  yellow,  and  a  precipitate  is  formed.  After 
standing  for  12  hours  filter  off  this  precipitate  and  wash  it  with 
water  until  the  last  washing,  on  being  treated  with  ammonium 
oxalate,  stirred,  and  allowed  to  stand  for  some  time,  shows  no 
precipitation.  To  purify  the  product  which  still  contains  some 
nitrate,  divide  it  first  into  portions  of  20  grams;  then  dissolve 
each  portion  in  40  cc.  of  hot  water  containing  a  little  acetic  acid, 
filter  the  solution  quickly  through  a  plaited  filter  and  immedi- 
ately precipitate  the  salt  from  the  filtrate  by  adding  a  solution  of 


180  METAL  AMMONIA  COMPOUNDS. 

40  g.  ammonium  chloride.  Cool  the  mixture  and  after  letting  it 
stand  24  hours  filter  off  the  yellow  crystals,  wash  them  with 
90%  alcohol  until  the  washings  barely,  give  a  test  for  chloride, 
then  with  pure  alcohol,  and  dry  the  product  in  a  desiccator.  The 
yield  is  variable. 

Reactions.  A  cold  solution  of  croceochloride  on  being  treated 
with  potassium  chromate  gives  a  yellow  precipitate  of  short, 
blunt  crystals;  with  potassium  pyrochromate  there  appears,  (but 
only  after  shaking  some  time)  a  deposit  of  short  leaflets  which  are 
frequently  grouped  together  in  the  form  of  stars  (microscope). 
On  being  treated  with  nitric  acid,  yellow,  feathery  crystals  of  the 
difficulty-soluble  nitrate  are  formed  (0.25  g.  of  the  latter  dissolve 
in  100  cc.  of  water  at  the  room  temperature). 

138.    Comparison    of    the    Isomeric    Dinitritotetramminecobaltic 

Salts. 

The  difference  between  the  flavo  and  croceo  salts,  in  both  of  which  the 
complex  has  the  same  empirical  formula,  [Co(NH3)4X2],  can  be  explained 
according  to  Werner's  theory  by  assuming  that  the  substituents  in  the  "inner 
sphere"  are  arranged  differently  in  space.  It  is  evident  that  if  the  con- 
stituents of  the  complex  are  placed  around  the  central  cobalt  atom  some- 
what as  the  corners  around  the  middle  point  of  an  octahedron,  there  are  two 
possible  arrangements;  either  the  two  constituents,  X,  which  are  different 
from  the  other  four,  are  arranged  opposite  to  one  another  and  are  connected 
by  the  hypothetical  axis  of  the  octahedron  (trans-position),  or  they  are  situ- 
ated side  by  side  and  are  joined  by  one  edge  of  the  octahedron  (cis-position). 
If  the  corners  of  the  octahedron  are  numbered  from  one  to  six,  as  shown  below, 
then  the  first  arrangement  may  be  designated  as  1  •  6  and  the  latter  as  1  •  2. 

i3      m/co     v 

•trn   /_ 


6  X  NH3 

1,  6:  trans-position  1,  2:  cis-position 

The  usefulness  of  the  above  conception  of  isomerism  is  shown,  among 
other  ways,  by  the  fact  that  it  does  not  allow  isomers  to  exist  of  the  com- 
plexes in  which  only  one  of  the  substituents  is  different  from  the  others; 
in  reality  no  such  isomers  have  been  observed.  On  the  other  hand,  it  has  not 
yet  been  positively  proved  which  series  of  compounds  corresponds  to  the 
trans  and  which  to  the  cis  arrangement,  (cf .  A.  Werner,  Lehrbuch  der  Stereo- 
chemie,  Jena,  1904). 


DINITRITOTETRAMMINECOBALTIC  SALTS.  181 

A.  Both  Salts  are  Tetrammine  Compounds.     Dissolve  some  of 
each  salt  separately  at  the  laboratory  temperature  in  concen- 
trated sulphuric  acid  and  allow  the  deep-violet  solutions  to  stand 
for  12  hours.     Then   add  to  each,  while  cooling  with  ice,  some- 
what more  than  an  equal  volume  of  concentrated  hydrochloric 
acid,   drop   by  drop,   until  the   further   addition   causes   merely 
a   slight  effervescence.     Two  days   later  decant  the  liquid  from 
the  small,  glistening,  green  crystals,  shake  the  latter  with  alcohol, 
drain    them    on  a  hardened   filter    and  wash    them  with    more 
alcohol.     Both  preparations  are  identical  and  consist  of  dichloro- 
tetramminecobaltic    acid  sulphate    (praseo    salt).     To   establish 
the  identity,  transform  each  portion  either  into  chloropentammine- 
cobaltic     chloride     (1)    or     into     chloroaquotetramminecobaltic 
chloride  (2); 

1.  Dissolve  small  portions  of  each  preparation  in  dilute  ammo- 
nia   by    heating    gently.     Add    concentrated    hydrochloric    acid 
little  by  little  to  the  purplish-red  solutions  and  then  heat  for  half 
an  hour  on  the  water-bath.     This  brings  about  an  almost  quanti- 
tative  separation   of   chloropentamminecobaltic    chloride  in  the 
form  of  minute  crystals  which  can  be  identified  by  the  reactions 
given  in  No.  131. 

2.  Dissolve   a   small   portion   of  each   of   the   praseosulphate 
preparations  in  water  by  heating  gently,  and  when  the  solutions 
have  turned  a  deep,  violet-red,  which  takes  place  after  standing 
for  some  time  in  the  cold  or  more  quickly  if  warm,  add  an  equal 
volume    of   alcohol    and   some    concentrated    hydrochloric    acid. 
Let  stand  for  12  hours  more  and  then  collect  the  precipitate  on 
a  filter.     The  chloroaquotetrammine  salt  is  very  similar  externally 
to  the  chloropurpureo  salt,  but  differs  from  it  in  the  following 
reactions;     mercuric    chloride    gives    well-formed    crystals    only 
after  standing  for  some   time;   potassium   chromate   and  pyro- 
chromate  likewise  yield  precipitates  only  after  long  standing,  or 
not  at  all. 

B.  Distinguishing  Reactions.     Boil  small  portions  of  the  salts, 
each  with  1   cc.  of  concentrated  hydrochloric  acid.     The  flavo 
salt   gives   a   green  precipitate   of  praseo   salt   and  the   solution 
acquires   a  bluish   color.     The   croceo  salt,   on  the  other  hand, 
gives   a   dull-red  precipitate  -of   chloronitritotetramminecobaltie 
chloride. 


182  METAL-AMMONIA    COMPOUNDS. 

To  prepare  somewhat  larger  amounts  of  these  last  substances, 
dissolve  0.5  g.  of  the  flavo  or  croceo  salt  in  10  cc.  of  concen- 
trated hydrochloric  acid.  After  12  hours  filter,  drain  and  wash 
the  precipitate,  first  with  hydrochloric  acid  diluted  with  an  equal 
volume  of  water,  and  finally  with  alcohol.  Yield,  0.4  g.  of  each. 
The  chloronitritotetramminecobaltic  salt  dissolves  readily  in  35  cc. 
of  cold  water  and  recrystallizes  upon  the  addition  of  70  cc.  of 
concentrated  hydrochloric  acid  to  the  filtrate.  The  praseo  salt, 
which  decomposes  very  easily,  can  be  recrystallized  by  dissolv- 
ing it  quickly  in  about  200  cc.  of  cold  water,  cooling  the  filtrate 
with  ice  and  salt,  and  slowly  introducing  70  cc.  of  concentrated 
hydrochloric  acid. 

139.    Trinitritotriamminecobalt,  [Co(NH3)3(NO2)3]. 

Dissolve  45  g.  of  crystallized  cobalt  chloride  in  125  c.c.  of 
water.  Dissolve  50  g.  of  ammonium  chloride  and  67  g.  of  sodium 
nitrite  without  heating  in  375  c.c.  of  water  and  add  this,  together 
with  250  c.c.  of  20  %  ammonia,  to  the  first  solution.  Conduct 
a  rapid  current  of  air  through  the  mixture  for  four  hours  (cf.  No. 
107).  Divide  the  thick  brown  solution  in  three  evaporating 
dishes  and  allow  it  to  stand  three  days  in  the  open  air  or  under 
the  hood.  Filter  off  the  crystals  which  separate  and  wash  them 
with  cold  water  until  the  washings  are  nearly  free  from  chloride. 
Yield,  about  35  g. 

Divide  the  crude  product  between  two  filters  and  extract  it 
with  1.0-1.25  liters  of  hot  water  containing  a  little  acetic  acid. 
On  cooling,  yellow  to  yellowish-brown  crystals  of  trinitrito- 
triamminecobalt  separate  from  the  clear  filtrate.  Yield, 
20-25  g. 

Solutions  of  this  substance  do  not  conduct  the  electric  current 
nor  do  they  give  precipitates  with  mercuric  chloride,  with  oxalates, 

or  with  chromates. 

• 

140.   Hexamminechromic  Nitrate,  [Cr(NH3)6](NO3)3  and  Chloro- 
pentamminechromic  Chloride,  [Cr(NH3)5Cl]Cl2. 

Anhydrous  chromic  chloride,  when  in  contact  with  liquid  ammonia,  takes 
up  partly  five  and  partly  six  molecules  of  ammonia.  The  hexammine  salt 
dissolves  readily  when  the  mixture  is  treated  with  cold  water,  and  from  the 
filtrate  the  nitrate  can  be  precipitated  by  the  addition  of  nitric  acid.  The 


HEXAMMINECHROMIC  NITRATE.  183 

chloropentammine  chloride,  on  the  other  hand,  does  not  dissolve  in  the  cold 
water,  but  can  be  purified  by  suitable  recrystallization.1 

Chloropentamminechromic  chloride  (chloropurpureochromic  chloride)  fur- 
nishes an  excellent  illustration  of  the  fact  that  the  characteristic  properties  of 
these  complex  compounds  are  often  less  influenced  by  the  nature  of  the  metal 
present  than  by  the  type  of  combination.  Chloropurpureochromic  chloride 
is  so  similar  to  chloropurpureocobaltic  chloride  that  it  is  necessary  to  decompose 
the  substances  completely  in  order  to  prove  that  the  metal  is  different. 

Generate  ammonia  by  heating  500  g.  of  concentrated  ammonia 
solution  on  a  Babo  funnel,  dry  the  gas  by  passing  it  through  two 
towers  filled  with  lime  and  through  a  U-tube  containing  sodium 
hydroxide,  then  condense  it  to  a  liquid  in  an  Erlenmeyer  flask 
containing  8  g.  of  finely  powdered,  anhydrous  chromic  chloride 
(Xo.  43).  In  order  to  exclude  carbon  dioxide  from  the  condens- 
ing flask,  this  must  be  closed  with  a  stopper  and  provided  with  an 
inlet  and  an  outlet  tube.  For  the  cooling  agent  use  a  mixture 
of  carbon-dioxide-snow  and  ether  (temperature  about  —  80°) 
placed  in  a  small  beaker,  which  in  turn  is  placed  within  a  larger 
beaker  so  that  an  inclosed  air  space  between  the  glass  walls 
shall  act  as  an  insulator! 

When  the  chromic  chloride  has  combined  with  the  liquid 
ammonia,  forming  a  reddish-brown  mass,  stop  the  operation  and 
allow  the  excess  of  ammonia  to  evaporate  from  an  open  dish. 
Triturate  the  residue  with  30  c.c.  of  ice-cold  water,  filter  and  wash 
the  undissolved  salt  with  a  little  cold  water  until  the  filtrate  runs 
through  of  a  reddish  color.  Add  concentrated  nitric  acid  to  the 
filtrate,  which  precipitates  luteo-chromic  nitrate.  Again  filter, 
dissolve  the  salt  in  a  little  warm  water  containing  a  few  drops  of 
nitric  acid,  and,  by  the  addition  of  more  nitric  acid,  precipitate 
out  the  crystalline  hexamminechromic  nitrate.  Yield,  about  7  g. 

The  residue  from  the  treatment  with  ice  water  consists  of  chloro- 
purpureochromic chloride.  Boil  it  in  a  beaker  with  concentrated 
hydrochloric  acid,  cool,  add  some  water,  collect  the  salt  on  a 
filter  and  wash  it  with  a  little  cold  water.  Then  dissolve  the 
salt  as  quickly  as  possible  at  50°  in  400-500  c.c.  of  water  contain- 
ing a  few  drops  of  sulphuric  acid.  Filter  the  solution  at  once 
through  a  large  plaited  filter  and'*add  an  equal  volume  of  con- 


1  For  a  method  of  preparing  chloropurpureochromic  chloride  without  the 
use  of  liquid  ammonia,  cf.  O.  Christe^nsen;  J.  pr.  Chem.  (2)  23,  54  (1881)  and 
S.  M.  Jorgensen,  ibid,  20,  105  (1879). 


184  METAL-AMMONIA    COMPOUNDS. 

centrated  hydrochloric  acid.  Allow  the  beautiful  red  crystals 
which  separate  to  stand  an  hour  in  contact  with  the  mother- 
liquor,  then  drain  them  and  wash  first  with  20  %  hydrochloric 
acid  (1  pt.  cone.  HC1  :  1  pt.  H2O),  then  with  alcohol,  and  dry  the 
product  in  a  desiccator.  Yield,  about  5  grams. 

Transformation  of  the  Luteo  Salt  into  Chloropurpureo  Salt. 
Mix  a  solution  of  hexamminechromic  nitrate,  dissolved  in  eleven 
times  its  weight  of  hot  water,  with  an  equal  volume  of  concen- 
trated hydrochloric  acid  and  boil  gently  for  from  thirty  minutes  to 
an  hour.  When  the  liquid  has  become  cold,  filter  off  the  precipi- 
tated chloropentammine  chromic  chloride  and  wash  it  as  above. 

141.  Tetrammineplatinous  Chloroplatinite,  [Pt(NH3)4]  PtCI4  (Green 

Salt  of  Magnus). 

This  salt,  which  was  first  prepared  by  Magnus,  and  according  to  Werner's 
system  of  nomenclature  should  be  called  tetrammineplatinous  chloroplatinite, 
can  be  precipitated  from  a  solution  of  chloroplatinous  acid  by  ammonia  or 
from  a  solution  of  tetrammineplatinous  chloride  by  chloroplatinous  acid. 

Reduce  a  hot  solution  of  2  g.  commercial  chloroplatinic  acid 
(H2PtCl6.2H2O)  in  7  c.c.  of  water  very  carefully  by  the  addition 
of  dilute  sulphurous  acid,  a  little  at  a  time  and  toward  the  last 
drop  by  drop;  wait  before  each  fresh  addition  until  the  odor  of 
sulphurous  acid  has  entirely  disappeared  (when  the  reaction  is 
nearly  complete  this  requires  some  time),  and  test  whether  a  drop 
of  the  reddish-yellow  solution  will  still  give  a  precipitate  when 
brought  in  contact  with  ammonium  chloride  on  a  watch  glass. 
An  excess  of  sulphurous  acid  would  decolorize  the  solution  and 
form  hydroplatinosulphurous  acid.1  When  a  precipitate  is  no 
longer  obtained  with  ammonium  chloride,  or  at  most  but  a  slight 
one,  heat  the  solution  of  chloroplatinous  acid  to  boiling  and  pre- 
cipitate the  salt  of  Magnus  in  the  form  of  dark  green  needles  by 
adding  an  excess  of  concentrated  ammonia.  The  amount  of 
precipitate  increases  considerably  as  the  solution  cools,  and  some- 
times, besides  the  above  salt,  which  settles  rapidly  to  the  bottom 

1  The  platinum  could  be  recovered  from  this  decolorized  solution  in  the 
form  of  the  salt  ammonium  platinosulphite  (NH4)2  [(SO3)2Pt].  H2O  by  making 
alkaline  with  ammonia,  evaporating  to  about  10  cc.  on  the  water-bath,  fil- 
tering, replacing  the  evaporated  ammonia  with  a  few  drops  of  concentrated 
ammonia  solution,  and  precipitating  with  alcohol.  Beautiful  white  needles 
(J.  Liebig,  1829.) 


ISOMERS  OF   DICHLORODIAMMINEPLATINUM.  185 

of  the  beaker,  a  little  of  a  yellowish,  crystalline  substance  is 
formed  which  probably  has  the  formula  [Pt(NH3)2Cl2];  this  last 
compound  does  not  settle  so  readily  and  can  be  decanted  off  with 
water.  Purify  the  product  by  draining  it  with  suction  and  wash- 
ing with  water,  alcohol  and  ether.  The  yield  is  slight  (about  0.25  g.  ) 
because  the  greater  part  of  the  platinum  has  remained  in  the 
mother-liquor  as  tetrammineplatinous  chloride.  Add  three  times 
its  volume  of  alcohol  to  this  solution,  filter  and  redissolve  the 
precipitate  (which  is  contaminated  with  ammonium  chloride)  in 
30  c.c.  of  hot  water.  Precipitate  Magnus'  salt  from  this  boiling 
solution  by  adding  chloroplatinous  acid  (prepared  by  reducing 
1  g.  of  chloroplatinic  acid  in  the  manner  described  above)  until  no 
more  green  precipitate  is  formed.  This  second  precipitation  of 
the  salt  is  usually  of  very  small  crystals  and  lighter  colored  than 
the  first.  The  total  yield  is  about  1.3  grams  (theoretical  1.7  grams). 

142.  Tetrammineplatinous  Chloride  [Pt(NH3)4]Cl2. 
This  salt,  which  was  formed  as  an  intermediate  product  in  the 
preceding  preparation,  may  be  obtained  pure  by  the  prolonged 
action  of  ammonia  upon  Magnus'  salt.  Treat  one  gram  of  the 
latter  with  40  c.c.  of  ammonia  in  a  small  flask,  and  boil  for  one  or 
two  hours  with  a  return-flow  condenser  until  finally,  after  the 
addition  of  20  c.c.  more  of  ammonia,  practically  all  of  the  salt  has 
dissolved.  Add  four  times  its  volume  of  alcohol  to  the  filtered 
solution,  concentrate  it  on  the  water-bath,  filter  off  the  brownish 
turbidity  which  forms,  and  allow  the  solution  to  crystallize. 
Free  the  crystals  from  mother-liquor  by  pressing  them  on  a  porous 
plate;  then  recrystallize  the  product  after  first  clarifying  the 
solution  with  a  little  bone-charcoal.  About  0.3  g.  of  pure  white 
crystals  are  obtained. 

143.     Isomers  of  Dichlorodiammineplatinum,  [Pt(NH3)2Cl2]. 

The  two  isomers  having  the  formula  [Pt(NH3)2Cl2]  (cf.  page  168)  have  been 
known  respectively  as  the  Chloride  of  the  Second  Base  of  Reiset  and  the  Salt 
of  Peyrone;  according  to  Blomstrand's  system  of  nomenclature  they  are 
platosammine  chloride  and  platosemidiammine  chloride;  whereas  Werner 
designates  them  as  1.3  or  Jrans-dichlorodiammineplatinum  and  1.2  or  cis- 
dichlorodiammineplatinum. 


Cl/       XNH3  NH,/ 

trans-dichlorodiamrnineplatinum  cis-clichlorodiammineplatinum 


186  HYDRATES. 

1.  Trans- Dichlorodiammineplatinum.      Heat   about   0.3   g.   of 
tetrammineplatinous  chloride  (cf.  No.  142)  to  250°  in  a  test-tube 
immersed  in  a  paraffin  bath.     Should  a  little  water  be  given  off 
at  first,  stop  the  experiment  temporarily  and  remove  the  drops, 
that  are  condensed  on  the  upper  walls  of  the  test-tube,  with  a 
narrow  piece  of  filter  paper.     Then  continue  the  heating  until 
ammonia  is  evolved;  the  substance  becomes  darker  colored,  and 
in  the  upper  part  of  the  tube  a  little  black  platinum  begins  to 
deposit  upon  the  glass.     The  treatment  requires  about  twenty 
minutes.     Extract  the  grayish-yellow  mass  in  a  beaker  with  20  c.c. 
of  boiling  water  and  filter  the  light-yellow  solution  while  hot  from 
the  black  residue;  on  cooling  an  entangled  mass  of  light-yellow 
needles  is  deposited.     Drain  the  crystals,  wash  with  alcohol  and 
ether  and  allow  them  to  dry.     Yield,  about  0.1  gram. 

2.  Cis- Dichlorodiammineplatinum.    Prepare  a  solution  of  chloro- 
platinous  acid  from  one  gram  of  commercial  chloroplatinic  acid, 
reducing  the  latter  with  sulphurous  acid  as  directed  in  No.  141. 
Concentrate,  the  solution  to  a  volume  of  about  2  c.c.  and  neu- 
tralize it  while  still  warm  with  a  concentrated  solution  of  ammo- 
nium carbonate.     Add  an  excess  of  the  latter  to  make  the  total 
volume    15   c.c.     Then   boil   the   solution,   whereupon   the   color 
changes  from  a  dark-reddish-brown  to  an  intense  yellow,  while  at 
the  same  time  green  crystals  of  Magnus'  salt  are  deposited.     Filter 
the  solution  while  it  is  boiling  hot,  cool  the  filtrate,  and  immedi- 
ately remove  the  yellow  crystals  that  separate  on  cooling.     Rinse 
these  crystals  with  alcohol  and  ether  and  recrystallize  the  product 
from  a  few  cubic  centimeters  of  boiling  water,  whereby  the  salt 
is  obtained  perfectly  free  from  the  green  Magnus'  Salt.     Yield, 
less  than  0.1  gram. 

Difference  between  the  Isomers.  The  first  isomer  is  of  a  light, 
straw-yellow  color,  and  the  mother-liquor  is  colorless.  The  second 
isomer  is  deeper  yellow,  it  is  more  soluble  in  water,  and  the  mother- 
liquor  is  yellow. 

HYDRATES. 

The  hydrates  formed  by  crystallization  from  aqueous  solutions  vary  in 
composition  with  the  temperature  and  concentration  of  the  solution.  A 
systematic  determination  of  the  hydrates  which  a  compound  can  form,  as 
well  as  of  the  conditions  under  which  they  can  exist,  may  be  carried  out  in 
the  light  of  the  principles  of  heterogeneous  equilibrium. 


HYDRATES. 


187 


1.  THERMIC  ANALYSIS.  If  the  freezing-point  of  an  aqueous  solution  of  a 
salt  is  determined  at  various  concentrations  and  the  results  are  plotted  with 
the  temperatures  as  ordinates  and  with  the  concentrations  in  per  cent  as 
abscissas,  a  curve  AC  is  obtained  which  starts  at  the  freezing-point  of  pure 
water  and  descends  in  accordance  with  the  law  of  Raoult  and  van't  Hoff 
(Fig.  23).  The  highly  concentrated  solutions,  on  the  other  hand,  are  to  be 
regarded  as  solutions  of  water  in  the  salt;  the  curve  CB  shows  the  region  in 
which  the  freezing-point  rises  as  the  amount  of  water  diminishes  until  finally 
the  freezing-point  (or  melting-point)  of  the  pure  salt  is  reached.  In  the 
region  of  the  first  curve,  that  part  of  the  mixture  which  acts  as  solvent  crys- 
tallizes out  on  freezing,  —  in  this  case  pure  ice;  in  the  region  of  the  second 
curve,  the  solid  salt  is  deposited  as  the  solution  cools.  The  point  where 
these  two  curves  intersect  is  called  the  eutectic  point,  and  is  characterized  by 
the  fact  that  salt  and  ice  crystallize  simultaneously  in  an  intimate,  eutectic 
mixture.  Such  mixtures,  which  are  also  known  as  cryohydrates,  possess 
freezing  (or  melting)  points  which  are  constant,  and  are  lower  than  the  freez- 
ing-point of  pure  water;  they  are  on  this  account  used  for  maintaining 
uniform  temperatures  of  below  zero  centigrade. 

If  the  salt  employed  forms  a  chemical  compound  with  water,  then  the  two 
systems  water/hydrate  on  the  one  hand  and  hydrate/anhydrous  salt  on  the 
other  hand  are  each  to  be  considered  independently,  according  to  the  princi- 
ple just  outlined.  If,  as  before,  the  concentrations  are  plotted  as  abscissas 
and  the  temperatures  as  ordinates,  a  double  pair  of  curves  is  obtained  with 
two  eutectic  points.  The  two  inside  curves  unite  in  a  common  maximum 
which  is  the  freezing-point  of  the  pure  hydrate  (Fig.  24).  Conversely  it  is 
clear  that  by  determining  the  freezing-points  of  a  complete  series  of  mixtures 


cone. 


FIG.  23. 


Compound  Cone. 

FIG.  24. 


of  the  salt  and  water,  the  presence  of  a  hydrate  can  be  detected  by  the  ap- 
pearance of  a  maximum  on  the  curve,  and  its  composition  can  be  read  by 
dropping  a  perpendicular  from  the  maximum  point  to  the  horizontal  axis. 
Several  maximum  points  indicate  the  presence  of  several  hydrates.  In  this  way 
Roozeboom  has,  for  example,  proved  in  the  case  of  ferric  chloride  the  exis- 
tence of  the  hydrates  2FeCl3.4H2O,  2FeCl3.5H2O,  2FeCl3.7H2O,  2FeCl3.12H2O. 
This  method  of  analysis  does  not  apply  solely  to  the  mixtures  of  a  salt 
and  water,  but  it  can  be  used  very  generally  to  prove  the  existence  of  com- 


188  HYDRATES. 

pounds.     It  has  acquired  a  high  importance  in  the  study  of  alloys  (Tammann 
and  his  students:   see  articles  in  Z.  anorg.  Chem.  mostly  later  than  1903). 

2.  VAPOR  TENSION  ANALYSIS.  When  a  hydrated  crystallized  salt  is  in 
equilibrium  with  water  vapor,  its  water  of  hydration  can  be  progressively 
withdrawn  in  the  same  manner  that  ammonia  is  withdrawn  from  the  metal 
ammonia  compounds  (cf.  p.  163).  Two  methods  are  available  by  which 
salts  can  be  investigated  from  this  point  of  view.  According  to  the  first 
the  vapor  pressure  of  the  salt  is  determined  for  varying  water  content,  and  a 
curve  is  constructed  with  the  aqueous  tensions  as  ordinates  and  the  corre- 
sponding percentages  of  water  as  abscissas;  if  the  salt  forms  a  number  of  dif- 
ferent hydrates,  the  vapor  pressure  above  any  given  mixture  of  the  hydrates 
remains  constant,  when  water  is  slowly  withdrawn,  as  long  as  any  of  the 
hydrate  richest  in  water  is  still  present.  When  this  hydrate  is  entirely 
exhausted,  the  pressure  sinks  abruptly,  and  a  second,  and  lower,  horizontal 
line  on  the  plot  corresponds  to  the  tension  of  the  next  lower  hydrate,  etc. 
(cf.  Fig.  25).  The  tension  of  the  aqueous  vapor  in  the  case  of  hydrated  cup- 
ric  sulphate  is  given  in  the  following  table: 

CuSO4  +  4.5  H2O 46.3  mm. 

CuSO4  +  3.5H20 47.1      " 

CuSO4  +  2.5H20 29.9     " 

CuSO4+  1.5H20 29.7     " 

CuSO4  +  0.5H20 4.4     " 

From  this  we  may  conclude  that  the  following  hydrates  exist:  CuSO4.5  H2O; 
CuSO4.3H2O;  CuSO4.lH2O.     The  vapor  pressure   of   the   pentahydrate    is 
about  46  mm.  and  remains  constant,   irrespective  of  the  amount  present, 
until  the  pentahydrate  has  entirely  disappeared.     Then  the  pressure  drops  to 
that  of  the  trihydrate  (about  30  mm.),  and  again  remains  constant  until 
this  is  completely  changed  into  monohydrate  (vapor  pressure  about  4. 5  mm.). 
The  other  method  of  vapor  tension  analysis  consists  in  determining  the 
decomposition   temperature   under   a   constant   pressure   of    aqueous   vapor 
(cf.  p.  163).     Approximate  values  for  the  decomposition  temperature  can 
be  obtained,  however,  by  finding  the  point  at  which 
water  is  given  up  when  the  salt  is  heated  in  any 
indifferent  atmosphere.      In  this  way  crystallized 
copper  sulphate  loses  the  last  molecule  of  water  at 
220°  to  240°,  and  from  this  and  the  fact  that  the 
other  molecules  of  water  escape  at  much  lower  tem- 
peratures, the  existence  of  a  definite  monohydrate 
has  been  recognized  for  a  long  time. 

In  the  cases  where  the  water  is  not  chemically 
bound  in  a  compound,  but  is  merely  adsorbed  (cf. 
pp.  34  and  38),  either  of  these  methods  of  investi- 
gation shows  merely  a  continuous  loss  of  water  with  rising  temperature,  or 
diminishing  pressure. 

The  question  as  to  how  the  water  is  united  with  the  components  of  the 


MELTING-POINT   MAXIMUM.  189 

salt  can  be  answered  only  in  a  few  special  cases.  Undoubtedly  in  certain 
metal-ammonia  compounds  it  belongs  to  the  complex,  positive  component  of 
the  salt,  (cf.  especially  No.  134).  Probably  the  water  is  also  attached  to  the 
metal  atom  as  one  of  the  constituents  of  the  complex  in  such  of  the  hydrated 
salts  as  are  closely  related  to  the  metal-ammonia  salts  in  their  stoichiomet- 
rical  composition.  For  example,  cupric  sulphate  hydrate  is  analogous  to 
tetramminecupric  sulphate  (No.  126),  whereas  the  violet  modification  of 
chromic  chloride  corresponds  to  the  hexammine  salts  (No.  140) : 

[Cr(NH3)6]Cl3 

[Cr(NH3)5H20]Cl3 

[Cr(H20)6]Cl3. 

The  fact  that  the  number  of  molecules  of  water  in  hydrated  salts  is  fre- 
quently four  or  six  is  quite  in  accord  with  the  views  just  stated,  inasmuch  as 
four  and  six  are  also  the  most  usual  coordination  numbers.  It  cannot  be 
claimed,  however,  that  all  compounds  containing  water  of  crystallization 
must  possess  complex  cations;  since,  among  other  reasons,  it  is  quite  certain 
that  the  negative  components  of  compounds  have  the  power  of  attaching 
water  molecules. 

How  much,  and  in  what  manner,  water  is  united  with  the  ions  in  solution 
is  not  positively  known.  A  number  of  facts  make  it  seem  extremely  prob- 
able, however,  that  the  ions  are  hydrated.1 

144.   The   Melting-point   Maximum   for   Magnesium   Nitrate 

Hexahydrate ;  Eutectic  Mixture  of  Barium  Chloride 

Dihydrate  and  Water. 

1.  The  freezing-point  diagram  of  a  series  of  mixtures  of  anhydrous  mag- 
nesium nitrate  and  water  shows  a  pronounced  maximum  for  the  composition 
corresponding  to  that  of  the  hexahydrate.  Measurements  from  which  such 
a  diagram  can  be  constructed  can  be  made  with  the  very  simplest  apparatus. 
Starting  with  the  crystallized  hexahydrate,  one  branch  of  the  curve  is 
obtained  through  successive  additions  of  water,  which  by  dissolving  in  the 
hexahydrate  depresses  its  freezing-point;  on  the  other  hand,  another  branch 
of  the  curve  is  obtained  by  adding  successive  portions  of  a  less  hydrated 
magnesium  nitrate  which  likewise  dissolves  in  the  hexahydrate  and  depresses 
its  freezing-point.  The  two  branches  of  the  curve  unite  in  a  very  pronounced 
maximum  at  the  composition  Mg(NO3)2  +  6H2O.  As  a  rule  the  maxima 
obtained  in  the  investigation  of  hydrates  are  less  pronounced,  and  often  they 
are  entirely  concealed.  Regarding  the  so-called  "concealed  maxima,"  see 
Tammann.  A  critical  compilation  of  all  the  known  equilibria  between 
water  and  inorganic  substances  is  given  in  Table  176,  in  Landolt-Bornstein- 
Meyerhoffer's  physikalisch-chemischen  Tabellen,  3d  edition,  1905. 

1  E.  W.  Washburn.  The  Hydration  of  Ions.  J.  Am.  Chem.  Soc.  31,  322 
(1909);  Hydrates  in  Solution;  Review  of  Recent  Experimental  and  Theoretical 
Contributions,  Technology  Quarterly,  1908,  360,  or  Jahrbuch  der  Radio- 
aktivitat  u.  Elektronik,  1908,  December. 


190  HYDRATES. 

The  method  employed  is  the  same  in  principle  as  that  followed 
in  determining  freezing-points  with  the  well-known  Beckmann 
apparatus,  but  it  is  very  much  more  simple  because  the  freezing- 
points  here  measured  are  so  far  apart  that  an  accuracy  of  between 
one  and  two  degrees  is  sufficient.  As  container  for  the  sub- 
stance, a  test-tube  about  19  cm.  long  and  2  to  2.5  cm.  wide  suffices; 
as  thermometer,  one  graduated  in  whole  degrees  and  reading 
to  120°  is  suitable;  an  air-mantle  is  superfluous;  as  stirrer,  a  thin 
glass  rod  or  a  piece  of  iron  wire  bent  into  a  suitable  shape  may  be 
employed,  although  it  is  preferable  to  use  a  heavy  platinum  wire 
as  in  the  Beckmann  apparatus.  As  heating  and  cooling  baths, 
two  beakers  of  from  300  to  400  c.c.  capacity,  filled  with  paraffin 
oil  and  provided  with  heavy  iron  wire  stirrers,  are  used.  Weigh- 
ings need  be  accurate  only  to  two,  or  at  the  most  three,  significant 
figures. 

Place  about  12  g.  of  magnesium  nitrate  hexahydrate  in  the  test 
tube  and  close  the  mouth  with  a  stopper  through  which  the 
thermometer  and  stirrer  are  inserted.  Dip  the  tube  well  into  the 
heating-bath,  heat  rapidly  to  70°,  and  from  that  point  on,  more 
slowly.  The  melting  point  of  the  substance  is  shown,  in  the 
first  place,  by  the  disappearance  of  crystals,  and,  second,  by  the 
fact  that  the  temperature  remains  for  some  time  constant  (89°). 
Next  heat  the  bath  to  110°-120°,  transfer  the  test-tube  to  the 
cooling-bath,  which  has  been  brought  to  about  90°,  and  observe 
the  point  at  which  crystals  begin  to  appear  and  the  temperature 
ceases  to  fall  during  the  crystallization;  by  repeating  the  experi- 
ment several  times  it  is  possible  to  establish  very  sharply  the  point 
at  which  the  crystals  appear  or  disappear.  Now,  from  a  weighed 
pipette,  or  from  a  small  weighed  wash-bottle,  add  one  or  two 
grams  of  water,  and  determine  the  amount  added  by  a  second 
weighing.  Mix  the  mass  well  with  the  water,  melt  it  completely, 
and  determine  as  before  the  point  where  crystals  begin  to  sepa- 
rate; it  lies  considerably  lower.  Follow  the  curve  in  this  way  by 
three  or  four  similar  experiments  until  the  laboratory  temperature 
is  reached.  The  heating-bath  should  be,  in  every  case,  about  10° 
warmer  and  the  cooling-bath  10°  colder  than  the  temperature 
which  is  to  be  measured,  the  latter  being  first  determined  by  a 
preliminary  rough  experiment.  As  the  concentration  of  the 
water  increases,  the  points  of  constant  temperature  become  less 


MELTING-POINT    MAXIMUM.  191 

pronounced,  but  the  beginning  of  crystallization  can  always  be 
observed  distinctly  by  the  clouding  of  the  liquid. 

To  establish  the  other  branch  of  the  curve,  determine  the 
freezing-points  in  a  new  series  of  mixtures:  First,  take  a  weighed 
amount  (about  30  g.)  of  the  hexahydrate,  and  free  it  from  a  part 
of  its  water  of  crystallization  by  heating  it  in  an  open  dish  upon 
the  water-bath  for  five  or  six  hours,  or  until  the  loss  in  weight 
corresponds  to  about  10%.  Pour  the  sirupy  liquid  while  still 
warm  into  a  porcelain  mortar,  and  stir  it  with  a  pestle  while  it 
solidifies.  When  cold,  powder  the  porcelain-like  mass  and  com- 
pute the  amount  of  water  present  by  igniting  a  weighed  sample 
until  it  is  completely  changed  to  MgO.  For  the  freezing-point 
determinations,  commence  as  in  the  first  series  with  the  pure 
hexahydrate,  and  add  successively  portions  of  1  — 3  g.  of  the  par- 
tially dehydrated  salt  just  prepared.  After  each  addition,  bring 
about  complete  solution  by  fusing  the  mixture,  and  then  observe 
the  point  at  which  crystals  begin  to  separate  on  cooling.  More 
highly  concentrated  liquids  are  sirupy  and  become  turbid  on 
account  of  air  bubbles,  and  thus  the  freezing-points  are  not  so 
sharply  defined;  it  is  sufficient,  however,  to  carry  the  observa- 
tions to  about  70°  on  this  branch  of  the  curve. 

To  construct  the  melting-point  diagram,  compute  for  each 
experiment  the  total  amount  of  anhydrous  magnesium  nitrate 
present,  the  total  amount  of  water,  and  from  these  the  number  of 
parts  of  anhydrous  salt  in  100  parts  of  the  mixture.  Plot  the 
latter  values  as  abscissas  and  the  observed  temperatures  as  ordi- 
nates.  Between  points  corresponding  to  42%  and  58%  of  the 
anhydrous  salt  the  curve  rises  from  about  18°  to  89°,  whereas 
between  58%  and  65^,  it  falls  from  89°  to  70°.  The  maximum 
point  of  the  curve  corresponds  to  the  hexahydrate,  and  it  is  thus 
proved  thermo-analytically  that  this  hydrate  exists  as  a  definite 
compound. 

2.  If  the  magnesium-nitrate-hexahydrate/water  diagram  were  to  be  carried 
out  to  the  freezing-point  of  pure  water,  it  would  show  the  existence  of  a  second 
hydrate  with  nine  molecules  of  water,  and  of  a  eutectic  point  at  — 29°,  between 
the  latter  and  pure  water.  The  •characteristics  of  the  eutectic  point  are  more 
conveniently  studied  in  the  case  of  the  system  barium-chloride-dihydrate  water. 

Place  10  c.c.  of  water  in  the  apparatus  used  above  and  determine 
the  freezing-point;  then  add  about  1  g.  of  crystallized  barium 


192  HYDRATES. 

chloride  (weighing  it  to  an  accuracy  of  two  significant  figures), 
and  determine  the  freezing-point  of  the  solution  with  a  thermom- 
eter which  reads  from  —  20°  to  100°  and  is  graduated  in  whole 
degrees.  Repeat  the  measurements  with  successive  additions  of 
1  g.  of  the  salt.  As  cooling-bath  a  mixture  of  ice  and  salt  can  be 
used  at  first.  The  solidification  point  is  easily  recognized  by  the 
temperature  remaining  constant,  but  this  point  is  usually  pre- 
ceded by  a  supercooling  of  about  0.5°  before  crystallization  is 
induced  by  vigorous  stirring.  When  temperature  equilibrium 
has  been  established,  withdraw  the  test-tube  from  the  freezing 
mixture  in  order  to  note  the  appearance  of  the  separated  solid. 
Pure  water  usually  shows  a  crust  of  ice  and  large  needles;  the 
solutions  show  finely  divided  flakes  of  ice.  If,  in  about  the  third 
experiment,  the  neighborhood  of  the  eutectic  point  is  reached,  a 
freezing-point  is  at  first  observed  as  usual,  but  after  the  mixture 
has  stood  in  the  freezing  bath  for  some  time  the  temperature 
again  falls  until  a  second  halting-point  is  registered  at  —  8.4°,  and 
at  this  temperature  opaque,  white  masses  of  the  cryohydrate,  con- 
sisting of  ice  and  barium-chloride-dihydrate,  separate.  In  the 
succeeding  experiments  the  solidifying  point  rises  rapidly  since 
the  region  is  reached  in  which  the  salt  acts  as  the  solvent  for  the 
water.  The  deposited  substance  now  consists  of  barium-chloride- 
dihydrate  which  is  easily  distinguishable  from  ice.  At  the  same 
time  the  halting-point  on  the  thermometer  becomes  less  distinct 
because  the  heat  of  solution  of  the  salt  is  less  than  the  heat  of 
fusion  of  ice.  It  is  possible,  however,  to  determine  very  sharply, 
as  in  the  magnesium  nitrate  series,  the  point  at  which  crystalli- 
zation begins;  and  this  is  especially  true  since  the  baths  used  in 
these  experiments  are  of  ice  water,  water  at  the  room  temper- 
ature, and  finally  liquid  paraffin,  instead  of  the  opaque  mixture  of 
ice  and  salt.  Each  measurement  is  repeated  several  times,  leav- 
ing a  little  of  the  barium  salt  undissolved  each  time  to  serve  in 
starting  the  next  crystallization.  It  is  sufficient  to  carry  the 
curve  up  to  about  70°  in  three  or  four  experiments.  The  tempera- 
ture readings  in  the  case  of  the  first  field  (the  ice  curve)  should  be 
accurate  to  about  0.1°,  and  in  the  second  field  (solubility  curve) 
to  within  one  or  two  whole  degrees. 

The  curve  is  plotted  as  before.     Between  0%  and  24%  BaCl2 
the  temperature  of  solidification  falls  from  0°  to  —  8.4°;  between 


CALCIUM  SULPHATE   HEMIHYDRATE.  193 

24%  and  about  33%,  it  rises  from  —  8.4°  to  above  70°.  If  the 
eutectic  point  itself  is  not  actually  obtained  in  one  of  the  experi- 
ments, it  is  easily  found  as  the  intersection  point  of  the  two  curves. 

145.    Calcium  Sulphate  Hemihydrate. 

Gypsum,  as  it  occurs  in  nature,  consists  of  calcium  sulphate  dihydrate. 
On  heating  gypsum,  its  aqueous  tension  increases  until  at  101.5°  it  is  equal  to 
the  atmospheric  pressure,  and  at  107°  it  is  970  mm.  or  the  same  as  that  of 
liquid  water  at  the  same  temperature;  in  a  closed  tube,  therefore,  water  and 
the  dihydrate  co-exist  up  to  a  temperature  of  107°.  At  higher  temperatures 
the  dihydrate  breaks  down  into  the  hemihydrate,  CaSO4.£H2O,  and  water, 
but  these  recombine  to  again  form  the  dihydrate  when  the  tube  is  allowed  to 
cool.  In  analogy  with  the  transformation  of  allotropic  forms  into  one  another, 
the  temperature  at  which  the  above  reaction  is  reversible  is  known  as  the 
transition  point. 

Gypsum  is  dehydrated  technically  by  heating  the  powdered  mineral  in 
iron  kettles  to  about  130°  while  stirring.  The  product  is  the  ordinary  plaster 
of  Paris,  which  consists  of  calcium  sulphate  hemihydrate,  and  which  when 
mixed  with  water  at  a  temperature  below  107°  takes  up  some  of  the  latter  to 
again  form  the  dihydrate.  This  solidifies  to  a  solid  mass  of  interlocking 
crystal  needles.  The  production  of  plaster  casts  is  accomplished  in  this 
manner.  When  heated  to  160°,  gypsum  becomes  "dead  burnt,"  or  com- 
pletely dehydrated,  in  which  form  it  does  not  readily  unite  with  water  again. 

The  conversion  of  the  dihydrate  to  hemihydrate  can  be  effected  below 
107°  if  the  aqueous  tension  is  lowered  by  the  presence  of  other  substances. 
Thus  under  a  saturated  solution  of  common  salt  the  hemihydrate  is  formed 
at  77°,  and  under  a  saturated  solution  of  magnesium  chloride  at  about  11°. 
The  dehydration  is  effected  very  rapidly  by  heating  gypsum  in  concentrated 
nitric  acid  (sp.  gr.  1.4)  on  the  water  bath. 

Prepare  calcium  sulphate  dihydrate  by  adding  120  c.c.  of 
2-normal  sulphuric  acid  to  an  aqueous  solution  of  50  g.  of  calcium 
chloride;  wash  the  precipitate  by  decantation  with  water,  and 
then  with  a  little  alcohol,  and  dry  it  in  the  hot  closet.  The  pro- 
duct consists  of  small  needles  (Microscope).  Mix  20  g.  of  the 
preparation  with  50  c.c.  of  concentrated  nitric  acid  (sp.  gr.  1.4) 
so  as  to  form  a  thick  paste,  and  heat  the  mixture  on  the 
water-bath  with  occasional  stirring.  Watch  the  course  of  the 
transformation  by  removing  a  few  drops  of  the  mixture  from 
time  to  time  and  examining  it  under  the  microscope;  after  about 
10  minutes  the  finely  pointed  needles  disappear,  and  compact 
prisms  with  right-angled  outlines  appear  in  their  place.  After 
half  an  hour  cool  the  contents  of  the  dish,  allow  the  solid  to  settle. 


194  HYDRATES. 

decant  off  the  liquid  as  completely  as  possible,  shake  the  residue 
with  50%  alcohol,  and  filter  at  once  with  suction.  Throw  away 
the  filtrate,  wash  the  precipitate,  and  dry  it  in  the  hot  closet. 
Yield,  almost  quantitative. 

Determine  the  water  content  of  the  product  by  igniting  a 
weighed  sample  to  faint  redness.  The  hemihydrate  contains 
6.2%  water. 

The  hemihydrate  prepared  as  above  consists  of  larger  crystals 
than  commercial  plaster  of  Paris,  and  for  this  reason  it  "  sets  " 
more  slowly.  Mix  half  of  the  preparation  with  water  to  form  a 
thick  paste,  and  allow  it  to  stand  until,  at  the  end  of  about  30 
minutes,  solidification  takes  place.  The  process  may  be  watched 
more  closely  by  placing  a  few  drops  of  the  fresh  mixture  under 
the  microscope;  after  about  20  minutes,  fine  needles  of  the  dihy- 
drate  are  seen  to  appear,  and  after  that  to  increase  rapidly  in 
quantity  while  the  compact  prisms  of  the  hemihydrate  disappear. 
Simultaneously  with  the  formation  of  the  crystal  needles  under 
the  microscope,  the  larger  sample  grows  hard.  This  change  may 
be  compared  with  the  similar  transformation  of  potassium  lead 
iodide  (cf.  No.  105).  Moisten  the  remainder  of  the  hemihydrate 
with  a  mixture  of  alcohol  and  water;  this  time  the  mass  solidifies 
much  more  slowly  because  the  vapor  pressure  of  the  water  is 
diminished  by  the  presence  of  the  alcohol. 

146.   Hydrates  of  Sodium  Sulphate.     Supersaturated  Solutions. 


The  solubility  of  ordinary,  crystallized  sodium  sulphate,  NajSC^  .  10  H2O, 
increases  rapidly  with  rise  in  temperature,  but  can  be  followed  only  to  32.4° 
because  at  this  temperature  the  solid  decahydrate,  standing  in  contact  with 
the  solution,  is  completely  dehydrated.  The  solubility  of  the  anhydrous 
salt  is  peculiar  in  that  it  diminishes  with  rise  of  temperature;  thus  at  32.4° 
the  saturated  solution  contains  the  maximum  amount  of  sodium  sulphate. 
Accordingly,  when  a  saturated  solution  is  allowed  to  evaporate,  the  anhy- 
drous salt  separates  if  the  temperature  is  above  32.4°,  whereas  the  decahydrate 
is  deposited  if  the  solution  is  below  this  temperature.  The  solutions,  how- 
ever, can  very  easily  become  supersaturated,  provided  they  are  protected 
from  dust,  etc.,  which  tends  to  start  crystallization.  By  allowing  such 
supersaturated  solutions  to  remain  in  the  cold,  a  different  hydrate, 
NajSCVTH-jO,  crystallizes  spontaneously,  but  since  this  salt  is  more  soluble 
than  the  decahydrate,  the  solution  standing  over  it  remains  supersaturated 
with  respect  to  the  decahydrate. 

Dissolve  170  g.  of  crystallized  sodium  sulphate  in  75  c.c.  of 


TRANSITION  POINT  OF   SODIUM   SULPHATE.  195 

water  at  32°,  filter,  and  divide  the  solution  equally  among  five 
small,  clean  flasks  which  have  been  freshly  rinsed  with  distilled 
water.  In  filling  the  flasks  take  care  not  to  wet  the  necks  with 
the  solution.  Immediately  stopper  each  flask  with  a  loose  plug 
of  cotton.  Allow  two  of  the  flasks  to  stand  in  the  ice-chest  over 
night,  or  longer,  until  the  heptahydrate  has  crystallized  out  in 
large,  closely  packed  crystals  which  fill  about  one-third  the  vol- 
ume of  the  liquid.  Allow  the  three  other  flasks  to  cool  to  room 
temperature;  a  thick,  oily  solution  is  obtained  from  which, 
although  it  is  supersaturated,  no  crystals  separate  even  when 
the  liquid  is  gently  rotated.  Inoculate  the  first  flask  with  a 
minute  fragment  of  sodium  sulphate,  whereupon  crystals  at  once 
begin  to  form  at  the  point  inoculated  and  grow  rapidly  until  the 
whole  contents  of  the  flask  have  solidified.  Place  a  trace  of 
sodium  sulphate  upon  a  piece  of  filter  paper,  then  brush  it 
off  as  completely  as  possible.  Tear  off  a  piece  of  this  filter 
paper  and  use  it  to  inoculate  the  contents  of  the  second  flask. 
Open  the  third  flask,  and  close  it  with  the  thumb,  which  has 
first  been  well  rinsed  with  distilled  water.  If  the  thumb  is  per- 
fectly clean  and  free  from  crystals  of  the  salt  and  particles  of 
dust,  the  solution  will  bear  shaking  without  the  appearance  of 
crystals. 

Introduce  a  trace  of  decahydrate  into  each  of  the  two  flasks 
which  contain  the  deposited  heptahydrate.  The  supernatant 
solution  then  crystallizes  into  the  decahydrate  so  that  the  two 
hydrates  are  obtained  together  in  the  flask  one  above  the  other. 
After  standing  a  long  time,  the  more  soluble  heptahydrate  goes 
over  into  the  less  soluble  decahydrate. 

147.   Transition   Point   of   Sodium   Sulphate, 
Na2SO4 .10H2O  <=±  Na2SO4  +  10  H2O. 

32.383°. 

If  a  mass  of  crystals  of  sodium  sulphate  decahydrate  is  heated  slowly,  the 
temperature  rises  steadily  until  32.383°  is  reached,  at  which  point  it  remains 
constant  for  a  considerable  time,  because  the  heat  then  received  from  the 
exterior  is  used  entirely  in  dehydrating  the  salt.  If  the  mixture  of  water 
and  sodium  sulphate  is  heated  to  a  higher  temperature  and  then  allowed  to 
cool  slowly,  the  temperature  again  remains  constant  at  the  same  point  as 
long  as  the  heat  of  hydration  is  sufficient  to  compensate  the  loss  of  heat  to 
the  surroundings.  Compare  the  corresponding  relations  in  the  change  in 


196  HYDRATES. 

state  of  aggregation  (No.  6,  p.  15).  Practical  advantage  may  be  taken  of 
this  behavior  of  hydrated  sodium  sulphate  for  producing  and  maintaining 
very  accurately  the  temperature  of  32.383°,  as,  for  example,  for  a  fixed  point 
in  thermometry.  By  placing  a  mixture  of  the  anhydrous  salt  and  the 
hydrate  in  surroundings  of  approximately  32  to  33°,  the  temperature  of  the 
mass  adjusts  itself  sharply  to  32.383°,  and  remains  constant  a  very  long 
time  at  this  point,  for  if  the  external  temperature  is  somewhat  higher,  the 
heat  conducted  into  the  mixture  is  absorbed  in  dehydrating  the  salt,  whereas 
if  the  surroundings  are  cooler  the  temperature  of  the  mixture  is  maintained 
by  the  heat  liberated  in  the  hydration  of  the  salt.  A  bath  composed  of  such 
a  mixture,  therefore,  can  be  used  as  a  very  delicate  thermostat.1 

The  transformation  MnCl2 . 4  H2O  <=±  MnCl2 .  2  H2O  +  2  H2O  takes  place  at 
58.089°  ±  0.005°.  Richards  and  Wrede,  Z.  phys.  Chem.  61,  313  (1907). 

Dissolve  100  g.  of  crystallized  sodium  sulphate  in  50  c.c.  of 
water  at  about  33°,  filter  the  solution,  and  cause  the  salt  to 
recrystallize  by  shaking  the  solution  and  cooling  it  under  the 
water  tap.  Drain  the  crystalline  meal  on  the  suction  filter,  wash 
the  salt  once  with  a  little  cold  water,  and  use  it  moist  in  the  follow- 
ing experiment.  Place  about  12  g.  of  the  crystals  in  a  test-tube 
and  insert  a  thermometer  which  is  graduated  in  tenths,  or  fifths, 
of  a  degree.  Place  this  test-tube  inside  a  slightly  larger  one  so 
that  an  air  space  of  about  2  mm.  separates  their  walls,  and  clamp 
the  two  tubes  thus  arranged  so  that  they  dip  into  a  beaker  con- 
taining about  600  c.c.  of  water.  Keep  the  temperature  of  the 
water  at  35°  to  36°. 

To  prepare  the  equilibrium  mixture  in  the  inner  test-tube, 
remove  this  from  its  air  mantle,  and  dip  it  directly  into  the  warm 
water  in  the  beaker.  When  its  contents  are  partly  melted,  wipe 
the  tube  and  replace  it  in  the  larger  tube.  The  transition  tem- 
perature is  quickly  established,  especially  if  the  mixture  is  stirred, 
and  remains  constant  for  more  than  half  an  hour.  By  removing 
the  inner  tube  and  allowing  it  to  cool  somewhat,  the  mixture 
again  assumes  the  transition  temperature.  Before  every  read- 
ing of  the  thermometer,  the  mass  should  be  stirred. 

The  anhydride  present  in  the  mixture  is  recognized  as  a  fine 
turbid  powder  in  the  presence  of  relatively  large  crystals  of  the 
decahydrate. 

1  Cf.  Richards,  Z.  phys.  Chem.  26,  690  (1898);  Richards  and  Marked,  43, 
465  (1903). 


ISOMERIC   CHROMIC   CHLORIDE   HYDRATES.  197 

148.    Isomeric  Chromic  Chloride  Hydrates,  CrCl3.6H2O. 

Hydrated  chromic  salts  possess  both  violet  and  green  modifications. 
Although  solutions  of  most  of  the  green  chromic  compounds  cannot  be  made 
to  yield  the  crystallized  salts,  the  chloride,  on  the  other  hand,  can  be  crystal- 
lized in  both  forms,  each  having  the  same  composition  CrCl3.6H2O.  The 
violet  salt  is  normal  in  its  behavior,  inasmuch  as  when  it  is  dissolved  in 
water  the  entire  chlorine  is  ionizable,  and  the  solution  thus  contains  three 
chlorine  ions  to  one  hydrated  chromic  ion.  The  formula  of  the  violet  modi- 
fication is  therefore  [Cr(H2O)6]Cl3.  On  the  other  hand,  the  solution  of  the 
green  salt  contains  chlorine  in  a  non-ionic  condition  —  according  to  Werner 
and  Gubser,1  two  of  the  chlorine  atoms  are  attached  in  the  complex.  Since, 
furthermore,  two  of  the  six  molecules  of  water  are  less  firmly  bound  in  the 
salt  than  the  other  four,  the  green  chloride  may  be  formulated: 

[Cr(H2O)4CyC1.2H2O. 

For  further  details,  consult  the  original  article  cited  in  which  the  above  ex- 
planation is  deduced  from  conductivity  measurements.  It  has  been  shown 
by  Weinland  and  Koch2  that  precipitations  with  silver  salts  serve  qualita- 
tively but  not  quantitatively  to  explain  the  relations. 

Preparation  of  the  Crude  Chloride.  Warm  100  g.  of  chromic 
acid  anhydride  (under  the  hood  in  the  hydrogen  sulphide  room) 
in  a  flask,  with  400  g.  of  concentrated  hydrochloric  acid.  Red 
vapors  are  first  evolved.  Boil  the  solution  until  it  becomes  pure 
green  and  no  more  chlorine  is  evolved.  About  three  hours  are 
required,  and  during  the  boiling  more  hydrochloric  acid  is  added 
if  necessary.  Concentrate  the  solution  in  an  evaporating  dish 
until  it  has  the  consistency  of  sirup,  allow  it  to  cool,  spread  the 
thick  mass  of  crystals  on  a  porous  plate,  and  finally  dry  the  pro- 
duct in  a  desiccator  over  lime. 

1.  Crystallized  Green  Chromic  Chloride.  Dissolve  50  g.  of  the 
impure  chloride  in  40  c.c.  of  water,  filter,  and  while  keeping  cold 
with  a  mixture  of  ice  and  salt,  saturate  the  solution  with  gaseous 
hydrogen  chloride.  After  standing  several  hours,  drain  the 
crystalline  paste  in  a  funnel  containing  a  marble  around  which 
a  thin  cord  of  asbestos  is  laid.  Without  washing,  dry  the  crys- 
tals one  to  two  days  in  a  desiccator;  then  stir  them  up  with 
acetone,  which  does  not  dissolve  the  dry  salt,  drain  the  product 
on  a  hardened  filter,  and  wash  it  with  acetone  until  the  latter  runs 
through  colorless.  Yield,  10  to  20  grams. 

1  Ber.  34,  1591  (1901). 

2  Z.  anorg.  Chem.  39,  296,  320  (1904). 


198  HYDRATES. 

The  preparation  dissolves  in  water  to  an  emerald-green  solu- 
tion. If  this  solution  is  precipitated  with  ammonia,  and  the 
chromic  hydroxide  is  redissolved  in  hydrochloric  acid  while  cool- 
ing, a  violet  solution  of  the  other  modification  of  the  chloride  is 
obtained.  The  violet  modification  is  the  more  stable  in  solu- 
tions that  are  free  from  acid. 

2.  Crystallized  Violet  Chromic  Chloride.  Dissolve  50  g.  of 
•crude  chromic  chloride  in  40  c.c.  of  water  in  a  flask  and  add  chro- 
mic hydroxide  hydrogel  to  the  boiling  solution  until  blue  litmus 
paper  held  in  the  escaping  vapors  is  no  longer  reddened.  The 
hydrogel  is  prepared  by  adding  ammonia  to  a  boiling  dilute  solu- 
tion of  from  10  to  20  g.  of  the  crude  chloride,  filtering  upon  a 
large  plaited  filter  and  washing  with  hot  water.  After  the  neu- 
tralized solution  has  been  boiled  for  half  an  hour,  whereby 
its  volume  is  reduced  to  about  50  c.c.,  cool  it  with  ice  and 
salt,  and  saturate  it  with  hydrogen  chloride  gas,  shaking  from 
time  to  time.  The  gas  should  be  introduced  slowly  at  first,  and 
the  temperature  of  the  solution  should  not  rise  above  0°.  Let 
the  very  fine-grained  precipitate  settle  for  several  hours,  then 
drain  the  supernatant  liquid  through  a  layer  of  asbestos  felt  on 
a  perforated  plate,  and  finally  suck  the  crystalline  paste  free 
from  liquid  on  the  filter.  Wash  with  a  little  very  cold  concen- 
trated hydrochloric  acid,  and  then  thoroughly  with  acetone;  the 
latter  washing  is  best  accomplished  by  removing  the  greater 
part  of  the  crystals  to  a  porcelain  dish,  stirring  them  with  ace- 
tone, and  again  draining  them  on  the  asbestos  filter. 

Dissolve  the  crude  product  thus  obtained  at  once  in  as  little 
cold  water  as  possible  (20  c.c.  at  the  most),  filter  the  solution, 
and,  while  keeping  it  very  cold,  saturate  it  with  hydrogen  chlo- 
ride gas.  The  violet  chloride  is  precipitated  almost  quantita- 
tively in  somewhat  larger  crystals  than  at  first.  After  some 
time  drain  the  precipitate  on  a  small  piece  of  linen  cloth  placed 
on  a  perforated  plate;  wash  the  product  with  acetone,  and  dry  it 
in  a  vacuum  desiccator  over  sulphuric  acid.  Yield,  5  to  14  g.1  of 
small,  violet  crystals  which  dissolve  in  water,  giving  a  violet  solu- 
tion; by  precipitating  this  solution  with  ammonia  and  dissolving 
the  precipitate  in  cold  hydrochloric  acid,  a  violet  solution  is  again 
obtained  (see  above). 

1  A  considerably  better  yield  can  be  obtained  from  a  solution  of  chromic 
nitrate  in  hydrochloric  acid.  N.  Bjerrum,  Z.  phys.  Chem.  59,  340  (1907). 


CHAPTER  VI. 


COMPLEX  NON-ELECTROLYTES. 

IN  this  chapter  are  brought  together  a  number  of  different  kinds  of  sub- 
stances, many  of  them  containing  organic  radicals,  and  all  of  them  showing 
very  little,  if  any,  tendency  to  undergo  electrolytic  dissociation.  The  ability 
of  these  substances  to  dissociate  otherwise  than  electrolytically  is  slight,  and 
most  of  them  can,  like  many  of  the  pure  organic  compounds,  be  distilled  with- 
out undergoing  decomposition.  It  is  characteristic  of  them,  however,  that  in 
the  presence  of  water  they  suffer  hydrolysis  (saponification)  instead  of  elec- 
trolytic dissociation  (cf.  p.  60).  Here,  as  in  the  case  of  organic  substances,  it 
is  permissible  to  develop  structural  formulas,  and  thus  the  probable  structure 
of  the  products  of  hydrolysis  can  be  derived.  For  example,  it  is  possible  to 
deduce  the  structure  of  certain  inorganic  acids  from  the  graphic  formulas  of 
their  chlorides  and  esters.  Cf.  Cyanic  Acid,  p.  113;  Oxy-acids  of  Sulphur, 
p.  126;  Acids  of  Phosphorus  p.  135. 

ACID  CHLORIDES. 

149.    Sulphuric  Acid  Dichloride   (Sulphuryl  Chloride),  SO2C12,  and 
Sulphuric  Acid  Monochloride  (Chlorosulphonic  Acid),  HO.SO2C1. 

By  the  term  acid  chloride  is  understood  a  substance  which  is  converted 
into  an  oxy-acid  when  its  chlorine  atoms  are  replaced  with  hydroxyl  groups. 
Sulphuryl  chloride,  SO2C12  is  the  chloride  of  sulphuric  acid: 

SO2C12  +  2  H2O  =  SO2(OH)2  +  2  HC1. 

It  is  produced  by  the  direct  union  of  sulphur  dioxide  and  chlorine  in  the  sun- 
light; the  combination  takes  place  more  readily,  however,  in  the  presence  of 
catalyzers  such  as  anhydrous  acetic  acid,  porous  charcoal,  or,  most  efficient 
of  all,  camphor. 

Sulphuric  acid  monochloride  is  formed  by  the  partial  hydrolysis  of  sul- 
phuryl  chloride,  but  it  can  also  be  prepared  from  sulphuric  acid  and  phos- 
phorus pentachloride  by  a  reaction  which  is  of  very  general  applicability. 
Phosphorus  pentachloride  acts  upon  substances  which  contain  hydroxyl  in 
such  a  way  that  phosphorus  oxychloride,  hydrogen  chloride,  and  a  chloro- 
substitution  product  of  the  original  material  are  formed: 
R.OH  +  PC15  =  RC1  +  POC13  +  HC1 

In  certain  cases  the  phosphorus  oxychloride  itself  also  acts  as  a  chlorinating 
agent. 

199 


200  ACID  CHLORIDES. 

Sulphuric  acid  monochloride  can  be  prepared  in  still  another  way  by  the 
direct  addition  of  hydrogen  chloride  to  sulphur  trioxide: 

SO3  +  HC1  =  HO.SO.C1. 

All  acid  chlorides  have  a  choking,  often  very  disagreeable  odor,  and  all 
fume  when  exposed  to  moist  air.  In  preparing  them,  moisture  must  be 
excluded  with  great  care. 

Sulphuryl  Chloride.  Connect  in  series  a  500-c.c.  distilling  flask, 
a  fairly  long  condenser,  and  a  receiving  flask,  making  all  the  joints 
tight  with  closely  fitting  cork  stoppers.  From  the  receiver  lead 
an  escape-tube  to  the  ventilating  flue.  Through  the  cork  in  the 
neck  of  the  distilling  flask  pass  two  tubes,  reaching  to  the  bottom 
of  the  flask,  by  means  of  which  sulphur  dioxide  and  chlorine  can 
be  introduced  separately.  Allow  the  distilling  flask  to  rest  in  a 
porcelain  dish  on  a  water-bath,  which  is  not  heated  at  the  start. 

Place  50  g.  of  camphor  in  the  distilling  flask  and  fill  the  porce- 
lain dish  wth  water  and  a  few  pieces  of  ice.  Generate  sulphur 
dioxide  from  400  g.  of  copper  turnings  and  800  g.  of  concentrated 
sulphuric  acid  (or  from  bisulphite  solution,  see  note,  p.  71). 
Pass  the  gas  first  through  a  sulphuric  acid  wash  bottle  and  then 
into  the  distilling  flask,  where  it  is  taken  up  by  the  camphor,  with 
which  it  forms  a  colorless  liquid.  Then  begin  to  introduce  chlorine, 
which  is  likewise  dried  by  sulphuric  acid  (cf.  No.  42,  p.  69). 
Regulate  the  evolution  of  the  two  gases  so  that  about  equal 
amounts  of  each  bubble  through  the  washing  bottles;  an  excess  of 
chlorine  colors  the  contents  of  the  distilling  flask  yellow.  Con- 
tinue the  process  until  this  flask  is  a  little  more  than  half  filled. 
Towards  the  end  allow  first  an  excess  of  chlorine  to  collect  in  the 
flask  and  then  strengthen  the  stream  of  sulphur  dioxide  until  this 
excess  is  just  removed,  after  which  stop  the  evolution  of  the  gases. 

After  some  time  —  six  to  twelve  hours  —  remove  the  porcelain 
dish  and  heat  the  flask  on  the  water-bath  as  long  as  anything 
distils  over;  at  first  a  considerable  quantity  of  gas  is  evolved.  If 
the  distillate  contains  free  chlorine  remove  it  by  shaking  the 
liquid  with  mercury  (cf.  Nos.  51  and  52),  and  filtering  through 
a  perfectly  dry  asbestos  felt  in  a  Gooch  crucible.  Finally  redistil 
the  material  from  a  distilling  flask  provided  with  a  thermometer 
and  a  condenser.  Boiling-point,  69.5°. 

A  mixture  of  a  few  drops  of  sulphuryl  chloride  and  a  few  c.c.  of 
water  reacts  slowly  with  the  formation  of  sulphuric  and  hydro- 


SULPHURIC   ACID   MONOCHLORIDE.  201 

chloric  acids.     Larger  quantities  react,  after  some  time,  suddenly 
and  very  energetically,  with  a  considerable  evolution  of  heat. 

Sulphuric  Acid  Monochloride. 

(a)  From  Sulphuric  Acid  and  Phosphorus  Pentachloride. 

The  chlorination  of  sulphuric  acid  by  phosphorus  pentachloride  takes 
place  according  to  the  equation 

SO2(OH)2  +  PC15  =  SO2C1.OH  +  POC13  +  HC1. 

The  phosphorus  oxychloride  formed  acts  likewise  as  a  chlorinating  agent  in 
this  case: 

211,80,  +  POC13  =  2SO2C1.OH  +  HPO3  +  HC1. 

First  prepare  pure  sulphuric  acid  "  monohydrate  "  by  adding 
fuming  sulphuric  acid  to  the  ordinary  concentrated  acid  until  the 
specific  gravity  at  exactly  15°  is  1.84. 

To  200  g.  of  this  acid  in  a  liter  flask  add  150  g.  of  phosphorus 
pentachloride  in  small  portions  from  a  glass  spatula.  The  mix- 
ture becomes  somewhat  heated,  and  large  quantities  of  hydrogen 
chloride  escape.  (The  operation  should  be  carried  out  under  the 
hood,  or,  better  still,  out-of-doors.)  When  all  the  phosphorus 
pentachloride  has  been  added,  heat  the  flask  on  a  Babo  funnel 
until  the  evolution  of  hydrogen  chloride  has  ceased.  Transfer 
the  liquid  product  to  a  distilling  flask,  or  a  tubulated  retort,  in 
which  a  thermometer  is  inserted  through  a  ring  of  asbestos  cord 
(not  through  a  cork).  Slip  a  glass  tube,  40  cm.  long  and  1  cm.  in 
diameter,  over  the  side  arm  of  the  distilling  flask,  or  the  neck 
of  the  retort  to  serve  as  an  air  condenser;  it  is  not  necessary  to 
make  the  joint  tight.  Distil  until  the  temperature  has  risen  to 
about  165°.  To  purify  the  crude  product,  redistil  it  from  a 
fractionating  flask  With  a  side  arm  condenser  (Fig.  7,  p.  6). 
Boiling-point,  153°.  Yield,  120  to  150  grams. 

(6)  From  Sulphur  Trioxide  and  Hydrogen  Chloride.  Melt  some 
commercial,  80%  fuming  sulphuric  acid,  which  can  be  obtained 
in  small  sealed  flasks,  by  placing  it  for  a  short  time  in  a  warm 
place.  Add  200  g.  of  this  acid  to  a  large,  gas-washing  flask  which 
has  ground  glass  joints;  cool  the  flask  and  contents  to  the  room 
temperature  and  pass  in  a  vigorous  stream  of  gaseous  hydrogen 
chloride.  As  soon  as  the  mixture  becomes  warm,  cool  it  by 
surrounding  the  flask  with  ice.  When  the  hydrogen  chloride 
ceases  to  be  absorbed,  transfer  the  liquid  to  a  distilling  flask; 


202  ACID  CHLORIDES. 

provide  an  air  condenser  as  in  (a)  and  fit  a  thermometer  in  the 
neck  of  the  flask  by  means  of  asbestos  cord.  On  distilling  the 
liquid  the  dissolved  hydrogen  chloride  escapes  first.  Save  the 
distillate  between  150°  and  165°  and  purify  it  by  redistilling. 
Yield,  about  170  g.  of  sulphuric  acid  monochloride,  boiling-point 
153°. 

Dependent  preparation:  Pyrosulphuric  Acid  Chloride  No.  150. 

150.  Pyrosulphuric    Acid    Chloride    (Disulphuryl    Chloride), 

Q  /  S02C1 

\  SO2C1 

The  action  of  phosphorus  pentachioride  on  sulphuric  acid  monochloride 
results  in  the  formation,  not  of  sulphuric  acid  dichloride,  but  of  pyrosul- 
phuric  acid  chloride,  with  the  splitting  out  of  a  molecule  of  water: 
C1S02OH         C1SCU 

;o  +  H2o. 

Cl  SO2OH        Cl  SO2  / 

For  the  purpose  of  withdrawing  the  water,  however,  phosphorus  pentoxide 
serves  better  than  the  pentachioride.  It  is  upon  this  mode  of  forming  the 
acid  chloride  that  the  customary  constitutional  formula  of  pyro-  or  di-sul- 
phuric  acid  is  based. 

Introduce  20  g.  of  phosphorus  pentoxide  and  then  30  g.  of 
sulphuric  acid  monochloride  into  a  small  retort,  and  close  the 
tubulus  with  asbestos.  Distil  slowly  without  a  condenser  and 
catch  the  distillate  in  a  flask  that  rests  in  a  bath  of  cold  water. 

Redistil  the  first  product  from  a  small  flask  with  a  side-arm  con- 
denser and  with  a  thermometer  made  tight  with  asbestos  cord. 
The  liquid  boils  sharply  at  146°.  Yield,  20  to  25  grams. 

151.  Sulphurous   Acid    Chloride    (Thionyl    Chloride)    SOCLj.1 

Provide  a  liter,  round-bottomed  flask  with  a  stopper  through 
which  a  delivery  tube  and  the  lower  end  of  a  return  condenser  are 
inserted,  the  latter  so  as  to  just  pass  through  the  stopper  and  the 
former  so  as  to  reach  to  the  bottom  of  the  flask.  Place  this 
apparatus  under  the  hood,  and  add  500  g.  of  phosphorus  penta- 
chioride (not  less)  to  the  flask.  Introduce  sulphur  dioxide 
(from  200  g.  copper  and  400  g.  concentrated  sulphuric  acid), 
purifying  it  by  bubbling  it  through  sulphuric  acid  and  then  pass- 
ing it  through  a  tube  filled  with  crystals  of  potassium  sulphate. 


1  Another  method  for  preparing  thionyl  chloride  depends  on  the  reac- 
tion: SO3  +  SC12  =  SO2  +  SOC12. 


SULPHUROUS   ACID   CHLORIDE.  203 

As  soon  as  all  the  phosphorus  pentachloride  is  dissolved,  stop  the 
flow  of  gas  at  once,  and  subject  the  resulting  mixture  of  thionyl- 
chloride  and  phosphorus  oxychloride  to  a  careful  fractional  dis- 
tillation. 

Through  the  cork  in  the  neck  of  a  round-bottomed  flask  fit  a 
fractionating  tower,  35  cm.  high,  containing  a  30-cm.  column  of 
coarse  glass  beads.  Connect  the  side  arm  of  the  tower  with  a 
condenser  and  insert  a  thermometer  through  a  cork  placed  in  the 
top  of  the  tower.  On  the  first  distillation  collect  four  fractions: 
(1)  all  that  distils  up  to  82°;  (2)  between  82°  and  92°;  (3)  between 
92°  and  105°;  (4)  between  105°  and  115°.  The  quantity  of  the 
fractions  varies  from  110  to  180  grams.  Distil  each  one  of  these 
portions  separately  in  the  same  apparatus  —  except  that  a 
smaller  flask  is  now  used  —  observing  the  following  plan  of  pro- 
cedure: pour  fraction  (1)  into  the  distilling  bulb  and  place  the  same 
flask  in  position  again  as  receiving  vessel.  Distil  until  the  tem- 
perature reaches  82°,  add  the  contents  of  receiving  flask  (2)  to 
the  bulb  and  distil  again.  Continue  to  collect  in  receiver  (1) 
until  the  temperature  again  reaches  82°,  then  exchange  this  flask 
for  the  now  empty  receiver  (2).  When  the  temperature  reaches 
92°  add  the  contents  of  receiver  (3)  and  continue  to  distil  into 
receiver  (2)  until  92°  is  again  reached;  then  exchange  this  flask 
for  receiver  (3).  At  105°  add  the  contents  of  receiver  (4)  and 
continue  to  distil  into  receiver  (3)  until  105°  is  once  more  reached; 
then  change  the  receivers  and  collect  the  distillate  in  (4)  until 
115°  is  reached.  It  will  be  found  that  the  middle  fractions  of 
the  second  series  of  distillations  are  smaller  than  in  the  first 
series.  Repeat  the  fractionation,  whereby  the  middle  fractions 
become  still  smaller  with  a  corresponding  increase  of  the  end 
fractions.  Finally  use  only  the  two  end  fractions;  distil  each  by 
itself  from  the  same  apparatus  after  it  has  been  cleaned,  and  reject 
a  small  amount  of  each  both  at  the  beginning  and  the  end  of  the 
distillation.  About  165  g.  of  thionyl  chloride  of  boiling-point 
77°  to  79°  and  about  180  g.  of  phosphorus  oxychloride  of  boiling- 
point  109°  to  111°  are  thus  obtained.  The  yield  of  thionylchlo- 
ride  amounts  to  55-60%  of  the  theoretical,  the  percentage  yield 
is  much  lower  when  smaller  quantities  are  prepared  because  of 
the  losses  incidental  to  the  fractionation:  these  losses  are  rela- 
tively less  in  preparing  larger  quantities.  The  preparation  is  not 


204  ACID    CHLORIDES. 

entirely  free  from  phosphorus  compounds;  it  is,  however,  admir- 
ably suited  for  use  as  a  chlorinating  agent  in  organic  chemistry. 
Its  odor  is  suffocating  and  offensive. 

A  few  drops  of  thionyl  chloride  added  to  a  little  water  react 
slowly,  —  more  rapidly  on  warming,  —  yielding  hydrochloric  and 
sulphurous  acids. 

To  test  for  the  presence  of  phosphorus,  add  some  nitric  acid  to 
the  solution  just  obtained,  evaporate  to  dry  ness  on  the  water- 
bath,  dissolve  the  residue  in  a  little  water,  and  test  with  ammonium 
molybdate. 

Dependent  preparation:    Symmetrical  Ethyl  Sulphite,  No.  155. 

152.   Nitrosylsulphuric  Acid,  HOSO2.NO2. 

Although  it  contains  no  chlorine,  nitrosylsulphuric  acid  can  be  regarded 
as  in  the  same  class  with  the  acid  chlorides,  because  its  characteristic  atom 
grouping,  —  SO2 .  NO2,  is  entirely  analogous  to  that  which  is  present  in  sul- 
phurylchloride,  or  in  sulphuric  acid  monochloride.  It  is,  like  the  two  latter 
compounds,  converted  by  hydrolysis  directly  into  sulphuric  acid. 

Nitrosylsulphuric  acid  can  be  formed  by  the  interaction  of  sulphur  dioxide 
and  nitric  acid: 

SO2  +  HNO3  =  HOSO2.NO2 
or  of  sulphuric  acid  and  nitrous  acid: 

2H2SO4  +  N2O3  =  H2O  +  2HOSO2.NO2 

Because  of  its  formation  in  the  lead  chambers  of  sulphuric  acid  plants 
when  insufficient  water  is  supplied,  this  compound  has  long  been  known  by 
the  name  of  "chamber  crystals  "  (cf.  No.  86). 

Place  100  g.  of  anhydrous  nitric  acid  (cf.  No.  34)  and  25  g.  of 
anhydrous  acetic  acid  in  an  Erlenmeyer  flask  and  surround  the 
flask  with  a  freezing  mixture.  For  introducing  sulphur  dioxide, 
insert  a  tube  1  cm.  in  diameter  through  the  cork.  The  top  of  this 
tube  is  closed  with  another  cork  and  just  below  this  stopper  a 
side  arm  is  provided  through  which  the  gas  is  to  enter.  Whenever 
the  lower  end  of  the  tube  becomes  clogged  with  crystals,  the 
stopper  may  be  removed  for  a  moment  and  the  obstruction  dis- 
lodged with  a  stirring  rod  (cf.  No.  46).  Generate  the  sulphur 
dioxide  from  200  g.  of  copper  turnings  and  400  g.  of  concentrated 
sulphuric  acid  (or  from  bisulphite  solution),  and  pass  it  through  a 
sulphuric  acid  wash  bottle. 

Pass  the  gas  rapidly  into  the  reaction  flask,  shake  the  mix- 
ture from  time  to  time,  and  take  particular  care  that  it  is  kept 


ESTERS.  205 

well  cooled,  since  otherwise  an  energetic,  sometimes  explosive 
oxidation  may  occur  according  to  the  equation 

S02  +  2  HN03  =  H2SO4  +  2  NO2. 

This  last  reaction  is  more  likely  to  take  place  if  the  nitric  acid 
has  not  been  diluted  with  acetic  acid  as  recommended  above. 

Drain  the  thick  crystalline  paste,  wash  it  with  a  little  cold,  glacial 
acetic  acid,  then  freely  with  carbon  tetrachloride,  and  dry  the 
product  in  a  vacuum  desiccator  over  sulphuric  acid.  The  crys- 
tals obtained  in  this  manner  can  be  preserved  in  a  well-stoppered 
flask  for  a  long  time  unchanged.  Yield,  about  80  to  90  g. 

ESTERS. 

Esters  are  derived  from  acids  by  the  replacement  of  the  acid  hydrogen 
atoms  by  hydrocarbon  radicals;  they  are  often  called  "etherial  salts,"  but  are 
to  be  distinguished  from  ordinary  salts  by  their  inability  to  dissociate  elec- 
trolytically. 

HN03  CH3N03 

Nitric  Acid.  Methyl  nitrate  (methylester  of  nitric  acid). 

Of  the  methods  of  forming  such  substances  the  following  three  will  be  con- 
sidered here: 

1.  From  an  acid  and  an  alcohol  with  elimination  of  water: 

C2H5OH    +    HNO3    =    H2O    +    C2H5NO3  (No.  153) 

Ethyl  Alcohol  Ethyl  Nitrate 

2.  From  an  acid  chloride  and  an  alcohol  with  elimination  of  hydrogen 
chloride: 

SOC12         +       2  C2H5OH    =    2  HC1    +    SO3(C2H5)2    (No.  155) 

Thionyl  Chloride  Ethyl  Alcohol  Diethyl  Sulphite 

3.  From  the  silver  salt  of  an  acid  and  an  alkyl  halide: 

Ag2S03        +        2  C2H5I    =    2  Agl    +    S03(C2H5)2     (No.  156) 

Silver  Sulphite  Ethyl  Iodide  Diethyl  Sulphite 

The  first  method  of  formation  is  reversible ;  the  opposed  reaction,  by  which 
the  ester  is  broken  down,  is  termed  saponification  and  is  in  its  nature  quite 
identical  with  the  hydrolysis  of  salts.  Since  the  saponification  of  esters 
takes  place  much  more  slowly  than  the  hydrolysis  of  salts,  it  is  admirably 
adapted  for  the  study  of  reaction  velocities  (see  text  books  on  Physical  Chem- 
istry) .  Inasmuch  as  the  two  opposed  reactions  —  esterification  and  saponi- 
fication —  often  lead  to  an  equilibrium  in  which  all  four  of  the  reacting 
substances  are  present  in  finite  determinable  concentrations,  this  so-called 
"ester  equilibrium"  is  used  with  great  success  in  demonstrating  the  law  of 


206  ESTP:RS. 

mass  action.  This  law  requires  that  at  a  constant  temperature  the  relation 
expressed  in  the  following  equation  shall  hold  true: 

[Ester]  •  [Water] 
[Acid]  •  [Alcohol] 

It  is  evident  from  this  equation  that  to  obtain  a  favorable  yield  of  ester  the 
concentration  of  the  water  should  be  kept  as  low  as  possible.  This  can  be 
accomplished  in  practice  by  adding  concentrated  sulphuric  acid  to  the  reac- 
tion mixture.  On  the  other  hand,  saponification  is  favored  by  the  removal 
of  the  acid,  and  this  can  be  brought  about  by  the  addition  of  a  base. 

153.    Ethyl  Nitrate,  C2H5ONO2. 

Ethyl  nitrate  is  formed  from  nitric  acid  and  ethyl  alcohol  according  to 
Method  1  described  above.  Since,  however,  the  nitric  acid  may  also  act  as 
an  oxidizing  agent  upon  the  alcohol,  and  since,  moreover,  the  oxidation  is  so 
vigorously  catalyzed  by  the  lower  oxides  of  nitrogen  that  the  mixture  some- 
times decomposes  explosively,  the  expedient  is  adopted  of  adding  urea,  which 
removes  the  reduction  products  of  nitric  acid.  Urea  reacts  with  nitrous  acid 
according  to  the  equation 

CO(NH2)2  +  2  HNO2  =  CO2  +  2  N2  +  3  H2O. 

Place  in  a  150-c.c.  distilling  flask  38  g.  of  absolute  alcohol,  6  g. 
of  urea,  and  60  g.  of  concentrated  nitric  acid  (sp.  gr.  1.4)  which 
has  previously  been  boiled  with  0.5  g.  of  urea.  Distil  the  mixture 
on  the  water-bath,  and  when  about  one-third  has  passed  over,  add 
slowly  from  a  dropping  funnel  a  further  mixture  of  100  g.  of  con- 
centrated nitric  acid  (which  has  likewise  been  boiled  with  a  gram 
of  urea  and  subsequently  cooled  to  the  room  temperature)  and 
75  g.  of  absolute  alcohol.  Continue  the  distillation  until  nothing 
more  passes  over.  Purify  the  ester  from  acid  and  alcohol  by 
shaking  it  three  times  with  water  in  a  separatory  funnel,  using  a 
double  volume  of  water  for  the  first  shaking.  Free  the  layer  of 
ester  as  completely  as  possible  from  drops  of  water  by  pouring 
it  into  a  dry  flask,  and  let  it  dry  by  standing  several  hours  in 
contact  with  granular  calcium  chloride.  Finally,  distil  the  prod- 
uct with  the  bulb  of  the  flask  immersed  in  a  water-bath.  Yield, 
about  40  g.  of  a  colorless  liquid  which  boils  sharply  at  87°.  On 
setting  fire  to  a  few  drops  of  the  ethyl  nitrate,  it  burns  with  a 
pale  flame. 

By  allowing  larger  amounts  of  the  above  mixture  to  drop  into 
the  reaction  flask,  the  ester  may  be  prepared  conveniently  in  con- 
siderable quantity. 


SYMMETRICAL  DIETHYL   SULPHITE.  207 

154.   Amyl  Nitrite,  C5Hn.ONO,  and  Methyl  Nitrite,  CH3.ONO. 

Amyl  nitrite  is  easily, formed  by  the  action  of  nitrous  acid  on  amyl  alcohol. 
Since  neither  amyl  nitrite  nor  amyl  alcohol  are  miscible  with  water,  the  prog- 
ress of  this  esterification  is  not  retarded  by  the  presence  of  water. 

To  50  g.  of  amyl  alcohol  and  300  g.  of  cold,  20%  sulphuric  acid 
in  a  750-c.c.  flask  add  60  g.  of  sodium  nitrite  in  small  portions 
while  shaking  the  flask  continuously  and  cooling  under  the  water 
tap.  Wait,  before  adding  each  fresh  portion,  until  the  reaction 
from  the  previous  addition  is  ended  and  the  yellowish-red  vapors 
have  disappeared  from  the  flask.  About  30  minutes  is  required 
for  this  operation.  After  waiting  an  hour  longer,  separate  the 
layers  in  a  separatory  funnel,  wash  the  oily  layer  twice  by  giving 
it  a  rotary  motion  with  water  (but  do  not  shake  it,  as  an  emulsion 
then  forms  which  will  separate  into  layers  only  on  long  standing), 
free  the  liquid  from  drops  of  water,  and  let  it  dry  by  standing 
overnight  with  lumps  of  fused  calcium  chloride.  Distil  the  ester 
from  a  fractionating  flask  with  a  side-arm  condenser.  Boiling- 
point,  97°.  Yield,  about  60  grams. 

Methyl   Nitrite. 

Methyl  alcohol  and  amyl  nitrite  interact  readily  with  the  formation  of 
methyl  nitrite  and  amyl  alcohol.  Amyl  nitrite  is  much  used  in  organic 
chemistry  as  a  reagent  in  preparing  nitroso  and  diazo  compounds. 

Mix  amyl  nitrite  with  about  one-third  its  weight  of  methyl 
alcohol.  After  a  short  time,  or  sooner  on  warming  gently,  methyl 
nitrite  (boiling-point,  —  12°)  begins  to  escape  as  a  gas.  By  using 
25  g.  of  amyl  nitrite  and  7  g.  of  methyl  alcohol,  several  cylinders 
full  of  gas  may  be  obtained.  When  set  on  fire  it  burns  with  a  pale 
flame. 

155.    Symmetrical  Diethyl  Sulphite,  SO(OC2H5)2. 

There  are  two  isomeric  substances  of  the  composition  (C2H5)2SO3,  both  of 
which  are  to  be  regarded  as  esters  of  sulphurous  acid.  One  of  them  is  pro- 
duced by  the  action  of  thionyl  chloride  on  ethyl  alcohol,  and  upon  being 
saponified  it  yields  sulphurous  acid  and  alcohol;  the  other  can  be  formed 
from  a  sulphite  and  a  halogen  alkyl,  or  by  the  esterification  of  ethyl  sulphonic 
acid,  C2H5SO2.OH.  Since  it  is  only  possible  for  thionyl  chloride  to  have 
the  structural  formula  O  :  SC12,  the  ester  formed  from  it  must  have  the  cor- 
responding symmetrical  structure 

/  Cl  .  OC2H 

O  :  S  '      +2  HOC2H6  =  2  HC1  +  O  :  S  ' 

\  OC2H6 


208  ESTERS. 

On  the  other  hand,  in  ethyl  sulphonic  acid,  since  this  can  be  formed  by 
the  oxidation  of  ethyl  mercaptan,  C2H5  .  SH,  it  must  be  true  that  the  ethyl 
is  bound  directly  to  the  sulphur  atom,  and  that  therefore  the  compound  has 
the  structural  formula  C2H5.SO2.  OH,  while  its  ester  has  the  corresponding 
unsymmetrical  structure 


As  regards  the  sulphites  and  free  sulphurous  acid  there  are,  therefore,  two 
possible  formulas.  The  fact  that  solid  sulphites  react  with  halogen  alkyls  to 
form  unsymmetrical  esters  suggests  that  the  sulphites  have  an  unsymmetrical 
structure.  There  is  no  positive  proof  as  to  whether  the  constitution  of  free 
sulphurous  acid  is  represented  by  the  formula 

/OH  /OH 

O  :  S  or  by  the  formula  O2  S 

\OH  \H 

It  is  perfectly  possible  that  both  forms  of  the  acid  may  exist  together  in  a 
state  of  mobile  equilibrium  (cf.  Cyanic  Acid,  No.  73).  Such  a  condition  is 
known  as  tautomerism  or  dynamic  isomerism  (cf.  A.  Findlay,  The  Phase 
Rule,  p.  193.) 

Insert  a  dropping  funnel  through  a  cork  in  the  mouth  of  a  75-c.c. 
flask  with  side-arm  condenser  (Fig.  7,  p.  6),  so  that  the  stem  of 
the  funnel  reaches  just  to  the  bulb  of  the  flask.  Place  40  g.  of 
thionyl  chloride  (one-third  mol)  in  the  bulb  and  surround  it  with 
a  mixture  of  ice  and  salt.  Then  allow  31  g.  of  alcohol  (which  has 
been  dehydrated  by  standing  two  hours  over  a  large  quantity  of 
quicklime  and  then  distilled),  to  drop  into  the  flask.  The  alcohol 
should  be  added  very  slowly  at  first,  as  the  reaction  is  violent  and 
considerable  heat  is  evolved.  The  operation  requires  thirty  to 
forty-five  minutes.  Finally  let  the  liquid  become  warmed  to  room 
temperature,  whereby  hydrogen  chloride  escapes  freely  (Hood), 
exchange  the  dropping  funnel  for  a  thermometer,  and  distil. 
At  first,  in  addition  to  hydrogen  chloride,  a  little  alcohol  passes 
over,  then  about  38  g.  of  ethyl  sulphite  distils  between  130°  and 
160°.  Redistil  the  crude  product  after  cleaning  and  drying  the 
apparatus.  About  35  grams  of  the  pure  ester  are  obtained,  boiling 
at  158°  to  158.5°. 

The  ester  is  a  colorless  oil  of  an  agreeable  odor,  and  is  some- 
what more  viscous  than  water.  It  is  not  decomposed  by  water 
even  on  boiling,  but  when  treated  with  caustic  soda,  it  is  saponi- 
fied with  the  formation  of  ethyl  alcohol  and  sodium  sulphite. 
Warm  a  few  drops  of  the  ester  with  2  c.c.  of  caustic  soda  and  a 


UNSYMMETRICAL  DIETHYL  SULPHITE.  209 

few  drops  of  alcohol,  drive  off  the  latter  by  boiling  a  short  time, 
cool,  and  acidify  the  solution  with  dilute  sulphuric  acid.  Sul- 
phur dioxide  escapes  and  can  be  recognized  by  its  odor. 

Saponify  a  second  portion  in  like  manner  by  adding  barium 
hydroxide  together  with  a  few  drops  of  alcohol;  boil  off  the 
alcohol  and  acidify  the  cooled  solution  with  nitric  acid.  At  first 
no  change  is  noticed,  but  on  boiling  a  cloudiness  appears  which 
is  caused  by  the  precipitation  of  barium  sulphate. 

156.   Unsymmetrical   Diethyl    Sulphite,    C2H5.SO2.OC2H5. 

To  prepare  silver  sulphite,  which  is  to  serve  as  the  starting 
material  for  this  preparation,  pass  a  stream  of  sulphur  dioxide 
into  a  solution  of  150  g.  silver  nitrate  in  500  c.c.  of  water.  Keep 
the  mixture  cooled  with  water,  and  continue  the  process  until  a 
small  filtered  portion  of  the  solution  no  longer  gives  a  precipitate 
with  hydrochloric  acid.  Drain  the  precipitated  silver  sulphite 
immediately  and  wash  it  successively  with  water,  alcohol,  and 
ether;  then  dry  it  over  sulphuric  acid  in  a  vacuum  desiccator. 

The  next  day  place  the  silver  sulphite  and  1.1  times  its  weight 
of  ethyl  iodide  in  a  flask,  provided  with  a  return  condenser,  and 
allow  the  mixture  to  stand  overnight.  The  top  of  the  condenser 
should  be  closed  with  a  calcium  chloride  tube  to  exclude  moisture; 
the  outer  jacket  should  be  filled  with  water,  but  a  constant  flow 
need  not  be  maintained.  After  24  hours,  add  150  to  200  c.c.  of 
thoroughly  dry  ether  (which  has  stood  for  several  days  in  con- 
tact with  sodium  wire),  and  boil  the  mixture  six  hours  on  the 
water-bath  with  return  condensation,  still  protecting  from  atmos- 
pheric moisture  with  the  calcium  chloride  tube.  Filter  the  liquid 
from  the  silver  iodide1  and  wash  the  latter  with  ether;  distil  the 
ether  from  the  filtrate,  using  a  tower  containing  glass  beads.  Then 
fractionate  the  liquid  in  a  smaller  flask,  using  a  tower  with  a  column 
of  glass  beads  13  cm.  high  and  1.5  cm.  wide  (cf.  No.  151).  Reject 
the  first  runnings  up  to  100°;  collect  two  fractions,  first  from  100° 
to  200°,  and  second  at  above  200°,  and  ref ractionate  these  portions 
repeatedly  according  to  the  process  described  in  No.  151.  The 

1  Recover  the  silver  from  the  silver  iodide  residue  by  reducing  it  with  a 
warm  solution  of  sodium  hydroxide  and  grape  sugar,  and  then  melting  the 
silver  powder  obtained  with  twice  its  weight  of  sodium  carbonate. 


210  ESTERS. 

pure,  unsymmetrical  diethyl  sulphite  boils  sharply  at  214°-215° 
(760  mm.).  Yield,  12  to  18  grams  of  a  colorless  oil,  quite  similar 
to  the  symmetrical  ester.  It  is  immiscible  with  water,  but  is  slowly 
saponified  by  it.  Alkalies  in  alcoholic  solution  cause  saponifica- 
tion  to  take  place  more  quickly. 

C2H5S02.OC2H5  +  H2O  =  C2H6S02.OH  +  C2H5OH. 

Allow  four  drops  of  the  ester  to  stand  a  few  minutes  with  an 
equal  amount  of  alcoholic,  potassium  hydroxide  solution.  The 
mixture  becomes  warm  and  a  solid  separates.  On  acidifying 
with  sulphuric  acid  and  boiling,  no  sulphur  dioxide  is  evolved. 
This  behavior  distinguishes  the  unsymmetrical  from  the  sym- 
metrical ester. 

157.   Triethyl  Phosphate,  PO(OC2H5)3. 

Add  14  g.  of  freshly-cut  sodium  to  150  c.c.  of  absolute  alcohol, 
and  after  the  reaction  moderates,  heat  the  mixture  on  the  water- 
bath  with  a  return  condenser  until  all  the  metal  has  dissolved. 
Cool  the  sodium  ethylate  solution  with  ice  and  salt,  then  add, 
drop  by  drop,  31  g.  of  phosphorus  oxychloride  (No.  46);  each  drop 
reacts  energetically  and  much  heat  is  evolved.  After  some  time, 
drain  the  solution  on  a  Buchner  funnel  from  the  precipitated 
sodium  chloride  and  rinse  the  residue  with  a  little  anhydrous  ether. 
The  sodium  chloride  is  so  finely  divided,  however,  that  it  cannot 
be  removed  completely.  Distil  the  alcohol  and  ether  from  the 
filtrate,  placing  the  bulb  of  the  flask  on  the  water-bath  and  wrap- 
ping the  upper  part  and  neck  with  a  towel.  Pour  the  residual 
liquid  through  a  smaller  filter  —  it  still  comes  through  somewhat 
cloudy  —  and  distil  it  from  a  fractionating  flask  with  a  side-arm 
condenser  (Fig.  7,  p.  6).  Collect  the  first  runnings  up  to  110° 
and  the  main  portion  above  110°;  some  sodium  chloride  is  left 
in  the  flask.  Fractionate  the  distillate  which  comes  over  above 
110°  twice  more  from  the  same  distilling  flask,  after  cleaning  it;  a 
pure  product  is  readily  obtained  which  boils  sharply  at  217°-2180. 
Yield,  20  to  25  g.  Triethyl  phosphate  is  a  colorless  oil  having  a 
typical  etherial  odor.  It  is  miscible  with  water  and  is  saponified 
by  it  to  diethyl  phosphoric  acid. 


METAL-ORGANIC  COMPOUNDS.  211 

158.   Tetraethyl  Silicate,  Si(OC2H5)4. 

To  about  20  g.  of  silicon  tetrachloride  (No.  51)  in  a  small  flask 
allow  1.1  times  its  weight  of  absolute  alcohol  to  flow  slowly  from  a 
dropping  funnel.  The  alcohol  must  have  been  boiled  with  quick- 
lime for  several  hours  immediately  beforehand  and  then  distilled 
out  of  contact  with  moisture.  During  the  operation  hydrogen 
chloride  escapes  in  quantity;  it  is  not  necessary  to  cool  the  mix- 
ture since  it  becomes  warmed  but  little. 

Purify  the  crude  product  by  distilling  it  from  a  small  round- 
bottomed  flask  and  through  a  fractionating  tower  containing  a 
13  cm.  column  of  glass  beads.  Use  an  air  condenser  consisting 
of  a  tube  40  cm.  long  and  1  to  1.5  cm.  in  diameter,  which  is  loosely 
slipped  over  the  side-arm  of  the  tower;  a  cork  connection  is  unneces- 
sary. Collect  two  fractions  during  the  first  distillation  —  the  first 
one,  which  is  large  in  quantity,  between  160°  and  175°,  and  the 
second,  which  is  smaller,  between  175°  and  185°.  On  the  second 
fractionation,  collect  a  small  amount  of  first  runnings  and  then 
portions  from  167°  to  170°  and  from-  170°  to  180°,  respectively. 
Finally,  after  cleaning  and  drying  the  apparatus,  distil  the 
fraction  boiling  between  167°  and  170°  again,  whereby  nearly 
all  goes  over  between  168°  and  169°.  Yield,  13-15  g. 

Small  amounts  of  higher  boiling  by-products  are  formed  and 
consist  of  esters  of  metasilicic  and  polysilicic  acids. 

Tetraethyl  silicate  is  a  colorless  liquid  having  a  typical  etherial 
odor;  it  is  immiscible  with  water,  but  it  dissolves  in  dilute  alcohol 
and  gradually  undergoes  hydrolysis  in  that  solution. 

METAL-ORGANIC  COMPOUNDS. 

The  metal  in  the  metal-organic  compounds  is  bound  directly  to  hydro- 
carbon radicals  or  to  carbon  monoxide:  Zn(C2H5)2,  zinc  ethyl  (No.  159); 
Pb(C6H5)4,  lead  tetraphenyl  (No.  160);  Ni(CO)4,  nickel  carbonyl  (No.  161). 
The  low  boiling-points  of  these  compounds,  which  would  scarcely  lead  one  to 
suspect  the  presence  of  a  metal,  permit  in  most  cases  a  ready  determination 
of  the  vapor  density  and  thus  of  the  molecular  weight.  In  this  respect,  as 
in  fact  in  almost  their  entire  chemical  behavior  (e.g.  the  ready  solubility  in 
organic  solvents),  these  compounds  show  themselves  to  be  closely  related  to 
the  purely  organic  compounds.  It  is,  therefore,  not  surprising  that  the 
theory  which  has  found  its  most  specific  application,  as  well  as  its  greatest 
success,  in  the  field  of  organic  chemistry  —  namely  the  valence  theory  in  its 


212  METAL-ORGANIC  COMPOUNDS. 

original  form  —  should  have  originated  in  the  discovery  of  these  very  metal- 
organic  compounds.  It  was  not  until  Frankland  (1853)  had  recognized  that 
the  saturation  capacity  of  the  metals  in  these  compounds  could  be  measured 
by  the  number  of  univalent  organic  radicals  which  are  joined  directly  to  the 
metal  atom  that  this  combining  capacity  began  to  be  regarded  as  a  funda- 
mental property  of  the  metals  —  as  well  as  of  all  other  elements  —  which 
governs  their  behavior  in  all  of  their  compounds. 


159.    Zinc  Ethyl.     Zn(C2H5)2. 

Zinc  and  ethyl  iodide  when  heated  together  combine  to  form  ethyl-zinc 

iodide: 

Zn  +  I.C2H5  =  C2H5.Zn.I, 

and  upon  stronger  heating  the  ethyl-zinc  iodide  is  transformed  into  zinc 
ethyl  and  zinc  iodide: 

2  C2H6.Zn.I  =  Zn(C2H5)2  +  ZnI2. 

The  formation  of  ethyl-zinc  iodide,  which  does  not  always  prove  success- 
ful when  zinc  alone  is  used,  can  be  accomplished  with  certainty  if  instead 
of  zinc  an  intimate  mixture  of  zinc  and  copper  (zinc-copper  couple)  is 
employed. 

Zinc  ethyl  takes  fire  spontaneously  when  it  comes  in  contact  with  the  air, 
and  it  is,  therefore,  a  dangerous  substance.  Above  all,  care  should  be  taken  to 
prevent  its  coming  in  contact  with  the  skin,  for  the  burns  produced  heal 
only  with  difficulty.  With  suitable  precautions,  however,  it  is  possible  to 
work  quite  safely  with  zinc  ethyl;  it  can,  for  example,  be  transferred  from 
one  vessel  to  another  without  danger  if  a  plentiful  supply  of  carbon  dioxide  is 
allowed  to  flow  over  the  openings  as  well  as  to  fill  the  interiors  of  the 
vessels. 

Place  an  intimate  mixture  of  100  g.  zinc  dust  and  12  g.  of  finely 
powdered  and  sifted  copper  oxide  in  a  glass  tube  of  2.5  cm.  diam- 
eter. Lay  the  tube  in  a  shallow  trough  of  asbestos  paper  over 
a  row  burner  and  pass  into  it  a  stream  of  dry  hydrogen.  After 
proving  the  purity  of  the  hydrogen,  heat  the  mixture  gently, 
allowing  the  flames  to  only  just  touch  the  asbestos.  During  the 
reduction  the  mass  becomes  lighter  in  color  and  swells  up  to  a 
considerable  extent;  on  this  account  the  tube  ought  not  to  be 
more  than  half  filled  with  the  mixture  at  the  outset.  At  the  end 
of  half  an  hour,  if  no  more  water  vapor  escapes  from  the  tube, 
extinguish  the  flame  and  allow  the  material  to  cool  in  an  atmos- 
phere of  hydrogen. 

Place  100  g.  of  this  zinc-copper  mixture  and  100  g.  of  ethyl 
iodide  in  a  200  to  300  c.c.  Erlenmeyer  flask  and  attach  a  return 


LEAD  TETRAPHENYL.  213 

condenser.  Heat  the  mixture  on  the  water-bath  until  a  thick, 
grayish  mass  is  formed  and  no  more  ethyl  iodide  condenses  and 
drips  back  (20  to  30  minutes).  Then  remove  the  condenser  and 
place  in  the  mouth  of  the  flask  a  cork  which  is  already  fitted  with 
a  gas  delivery  tube  to  reach  into  the  lower  third  of  the  flask  and 
with  another  tube  leading  to  a  condenser.  Place  at  the  lower 
end  of  the  condenser,  to  serve  as  a  receiving  vessel,  a  100-cc. 
distilling  flask  with  a  side-arm  condenser  (Fig.  7,  p.  6).  All 
the  joints  should  be  made  air-tight  with  corks.  After  filling 
the  entire  apparatus  with  carbon  dioxide,  heat  the  Erlenmeyer 
flask  to  180  to  220°  in  an  oil-bath,  whereupon  zinc  ethyl  distils 
over. 

Some  time  after  the  distillation  is  finished,  disconnect  the  receiv- 
ing vessel  and  stopper  it  immediately  with  a  cork,  already  made 
ready,  which  carries  a  thermometer  and  a  carbon  dioxide  de- 
livery tube.  Keeping  a  plentiful  supply  of  carbon  dioxide  flowing 
through  the  apparatus,,  redistil  the  liquid.  Catch  the  first  run- 
nings of  zinc  ethyl  mixed  with  ethyl  iodide  in  a  test-tube  filled 
with  carbon  dioxide.  At  110°  interrupt  the  distillation,  and 
replace  the  test-tube  with  a  thick-walled  tube  drawn  out  near  its 
upper  end  preparatory  to  being  sealed  off.  Fill  this  tube  like- 
wise with  carbon  dioxide  and  distil  over  all  the  liquid,  heating 
the  entire  distilling  flask  at  the  last  by  fanning  it  with  the  flame. 
Then  seal  the  tube  immediately  with  the  blast  lamp.  Carefully 
avoid  any  access  of  air  at  this  or  any  previous  part  of  the  opera- 
tion, since  this  would  cause  immediate  ignition  of  the  zinc  ethyl. 
Boiling-point,  118°.  Yield,  about  30  g. 

160.   Lead  Tetraphenyl,  Pb(C6H5)4,  by  Means  of  Grignard's 
Reagent;  Diphenyl  Lead  Iodide,  Pb(C6H5)2I2. 

The  action  of  metallic  magnesium  upon  etherial  solutions  of  alkyl  halides 
results,  as  was  discovered  by  Grignard,  in  the  formation  of  compounds  in 
which  the  alkyl  radical  is  bound  directly  to  the  magnesium: 

C6H6Br  +  Mg  =  C6H5MgBr. 

The  compounds  formed  are  soluble  in  ether,  and  themselves  contain  ether; 
their  etherial  solutions,  which  constitute  the  so-called  "Grignard's  reagent," 
can  be  employed  for  transferring  alkyl  groups  to  different  metals  as  well  as  to 
organic  radicals  of  the  most  varied  types. 

2  PbCl2  +  4  C6H5MgBr  =  Pb  +  Pb(C6H5)4  +  2  MgCl2  +  2  MgBr?. 


214  METAL-ORGANIC  COMPOUNDS. 

Lead  Tetraphenyl.  Place  70  g.  of  perfectly  dry  ether,  which 
has  stood  several  days  in  contact  with  sodium  wire,  and  25  g.  of 
brombenzene  in  a  small  flask.  Add  3.7  g.  of  magnesium  ribbon 
which  has  been  scraped  clean  with  a  knife,  close  the  flask  with  a 
calcium  chloride  tube  and  allow  it  to  stand  24  hours  in  a  dish  of 
water.  When  all  the  magnesium  has  dissolved,  add,  while 
shaking,  24  g.  of  dry  powdered  lead  chloride  in  small  portions, 
and  allow  the  mixture  to  stand  for  two  days  with  occasional 
shaking.  Then  add  this  mixture,  a  little  at  a  time,  to  200  c.c.  of 
water  and  acidify  faintly  with  dilute  hydrochloric  acid.  Collect 
the  precipitate  (which  is  colored  dark  by  precipitated  lead)  on  a 
suction  filter,  wash  it  with  water,  and  dry  it  in  the  hot  closet. 
The  mass  becomes  lighter  colored  on  drying  and  weighs  about 
15  g.  Boil  the  material  with  two  successive  portions  of  100  c.c. 
each  of  benzene,  using  a  reflux  condenser,  and  concentrate  the 
combined  filtrates  in  a  distilling  apparatus1  to  about  75  c.c.  Color- 
less, glistening  prisms,  melting-point  222°  to  224°,  crystallize  from 
the  liquid.  Yield,  8  to  9  g. 

Further  concentration  of  the  mother-liquid  furnishes  but  little 
additional  product. 

Diphenyl-lead-iodide.     (CQHb)2  PbI2. 

Two  of  the  phenyl  groups  in  lead  tetraphenyl  can  be  replaced  with 
hydroxyl  groups,  or  with  acid  radicals: 

Pb(C6H5)4  +  2  I2  =  (C6H5)2PbI2  +  2  C6H5I. 

Dissolve  3  g.  of  lead  tetraphenyl  by  warming  it  with  60  g.  of 
chloroform.  After  cooling  add  carefully  a  cold  solution  of  3  g. 
iodine  in  a  little  carbon  disulphide  until  the  color  of  the  iodine 
just  fails  to  disappear.  Allow  the  yellow  solution  to  evaporate 
in  a  warm  place;  extract  the  residue  first  with  10  c.c.  and  then 
with  5  cc.  of  carbon  disulphide;  concentrate2  the  filtered  extract 
to  a  volume  of  between  5  and  10  c.c.;  and  allow  it  to  crystal- 
lize, after  the  addition  of  2  to  3  c.c.  of  absolute  alcohol.  Drain 


1  Benzene  vapors  are  inflammable  and  burn  with  a  smoky  flame. 

*  Small  quantities  of  carbon  bisulphide  and  other  inflammable  liquids  can 
be  evaporated  over  a  free  flame  in  an  open  vessel  if  the  vapors  are  drawn 
rapidly  away  through  a  tube  connected  with  the  suction  pump.  The  tube 
should  be  inserted  about  one-third  the  way  into  the  beaker  or  flask. 


NICKEL  CARBONYL.  215 

the  deep  lemon-yellow  crystals,  wash  them  with  a  little  alcohol, 
and  dry  them  at  a  gentle  heat.  The  product  melts  with  decom- 
position1 at  105  to  107°.  The  yield  is  not  quite  quantitative,  as 
some  lead  iodide  is  always  formed. 

Nickel  Carbonyl,2  Ni(CO)4. 

When  carbon  monoxide  is  conducted  over  very  finely  divided  nickel, 
four  molecules  of  the  gas  combine  below  100°  with  one  atom  of  nickel  to  form 
nickel  carbonyl  (boiling-point,  43°).  At  higher  temperatures  this  compound 
dissociates  and  the  nickel  separates  in  the  form  of  a  dust,  or  deposits  as  a 
mirror  on  the  walls  of  the  vessel.  In  making  the  preparation  a  careful  main- 
tenance of  the  proper  temperature  and  a  state  of  very  fine  subdivision  of  the 
metal  are  important;  nickel  is  obtained  in  the  desired  condition  by  reducing 
nickel  oxalate  in  a  current  of  hydrogen. 

Treat  a  hot  solution  of  nickel  sulphate  with  oxalic  acid,  then 
add  ammonia,  but  without  entirely  neutralizing  the  solution. 
The  precipitation  of  nickel  oxalate  in  this  way  is  not  quite 
quantitative,  but  the  product  can  be  washed  readily  by 
decanting  with  hot  water.  By  draining  the  precipitate  and 
drying  it  in  the  hot  closet,  a  light-green,  loose  powder  is 
obtained. 

The  carbon  monoxide  may  be  prepared  by  heating  30  g.  of 
crystallized  oxalic  acid  with  100  cc.  of  concentrated  sulphuric 
acid,  and  passing  the  gas  through  two  wash-bottles  containing 
concentrated  sodium  hydroxide  solution.  Collect  the  carbon 
monoxide  in  a  gasometer. 

The  reduction  of  the  nickel  oxalate  and  the  synthesis  of  the 
nickel  carbonyl  may  be  accomplished  in  the  same  apparatus. 
Connect  two  sulphuric  acid  wash  bottles  in  series  at  one  end  of  a 
combustion  tube  about  26  cm.  long,  and  beyond  them  place  a 
three-way  cock  through  which  either  hydrogen  from  a  Kipp  gen- 
erator or  carbon  monoxide  from  the  gasometer  may  enter.  The 
carbon  monoxide,  before  entering  the  tube,  must  be  further  purified 
from  traces  of  carbon  dioxide  by  passing  through  two  bottles 
containing  caustic  soda  solution.  Place  an  asbestos  heating 

1  Diphenyl    is    formed:      (C6H6)2PbI3  =  PbI2  +  C6H8.C6H8.      The    same 
decomposition  takes  place  slowly  even  in  a  solution  of  diphenyl-lead-iodide. 
The  diphenyl  is  easily  recognized  by  its  characteristic  odor. 

2  Mond,  Langer,  and  Quincke,  Chem.  News,  62,  97  (1890). 


216  METAL-ORGANIC  COMPOUNDS. 

box1  (cf.  Fig.  4,  p.  3)  around  the  combustion  tube  and  insert 
a  thermometer  into  the  chamber.  Heat  the  tube  by  means  of 
a  Bunsen  burner  with  a  flame  spreader,  but  do  not  place  the 
flame  immediately  under  the  lower  opening  of  the  box.  Clamp 
the  combustion  tube  so  that  it  slants  downward  a  little  and  pro- 
vide an  extra  burner  to  expel  any  water  which  may  condense 
inside  the  tube. 

Place  5  to  7  g.  of  nickel  oxalate  in  the  tube  between  two  loose 
plugs  of  asbestos,  replace  the  air  completely  with  hydrogen,  and 
heat  the  charge  to  300°  in  a  current  of  hydrogen  until  it  has 
become  black  and  no  more  water  vapor  escapes.  Avoid  a  higher 
temperature,  as  the  nickel  would  then  lose  its  condition  of  fine 
subdivision.  Then  allow  the  temperature  to  sink  to  between  80° 
and  100°,  and  replace  the  hydrogen  by  means  of  a  current  of  air- 
free  carbon  monoxide  (the  wash  bottles  and  connections  should 
be  filled  with  carbon  monoxide  before  beginning  the  operation). 
Test  the  escaping  gases  for  nickel  carbonyl  as  follows : 

1.  Insert  into  the  end  of  the  combustion  tube  a  cork  carrying 
a  tube  drawn  out  to  a  capillary  jet;  set  fire  to  the  escaping  gas; 
it  burns  with  a  brilliantly  luminous  flame  from  which  metallic 
nickel  is  deposited  upon  a  cold  piece  of  porcelain  in  a  form  resem- 
bling soot. 

2.  Connect  the  combustion  tube  by  means  of  rubber  tubing 
with  a  carefully  cleaned  glass  tube,  and  heat  the  latter  gently 
with  a  Bunsen  burner;    a  mirror  of  nickel  is  deposited  in  the 
heated  part.'    If  the  tube  is  heated  to  redness  the  nickel  is  then 
precipitated  in  the  form  of  a  powder. 

Nickel  carbonyl  may  be  condensed  by  passing  the  gas  into  a 
vessel  surrounded  by  a  freezing  mixture.  It  is  a  colorless  liquid 
which  boils  at  43°  under  751  mm.  pressure,  and  solidifies  at  —  25° 
to  a  mass  of  needle-shaped  crystals. 


1  The  openings  in  the  cover  of  the  heating  box  should  be  covered  with 
discs  of  asbestos  to  avoid  drafts. 


CHAPTER  VII. 


PREPARATION    OF   COMPOUNDS  OF   THE   RARE 
ELEMENTS   FROM   THEIR   MINERALS. 

IN  this  chapter  methods  are  described  for  preparing  compounds  of  some 
of  the  rarer  elements.  It  has  seemed  advisable  to  devote  to  these  elements  a 
special  chapter  in  which  the  chief  stress  is  laid  upon  the  methods  of  working 
up  the  natural  raw  materials,  and  upon  the  characterization  of  the  individual 
elements  by  the  aid  of  their  compounds,  rather  than  upon  a  classification  of 
the  compounds  according  to  types. 

162.   Lithium  Carbonate  from  Lepidolite,  Petalite,  or  Spodumene; 
Spectroscopic  Tests  for  Rubidium  and  Other  Metals. 

Mix  100  g.  of  the  finely  powdered  mineral  intimately  with  100  g. 
of  ammonium  chloride  and  200  g.  of  finely  powdered  calcium 
carbonate;  heat  the  mixture  in  a  clay  crucible,  at  first  gently  for 
half  an  hour,  and  then  strongly  for  an  hour  in  the  furnace. 
Break  up  the  sintered  mass  and  extract  it  about  ten  times  by 
boiling  with  750  to  1000  c.c.  of  water  in  a  large  evaporating  dish. 
Four  off  the  solution  each  time  through  a  plaited  filter  and  begin 
at  once  to  concentrate  the  filtrate  in  a  second  evaporating  dish. 
After  the  entire  filtrate  has  been  reduced  to  about  one  liter, 
make  it  alkaline  with  ammonia  and  add  ammonium  carbonate 
until  a  little  of  the  liquid  when  filtered  gives  no  further  precipi- 
tation with  the  reagent.  Drain  the  solution,  on  a  suction  filter, 
from  the  precipitated  calcium  carbonate,  and,  as  the  latter  con- 
tains some  lithium  carbonate,  dissolve  it  in  acetic  acid,  dilute  to 
about  1.5  liters  and  precipitate  this  solution  cold  with  oxalic 
acid.  After  the  precipitate  has  settled,  filter  and  add  the  filtrate 
to  the  main  solution  of  the  lithium  salt.  Evaporate  the  entire 
solution  to  dryness,  and  drive  off  the  ammonium  salts  from  the 
residue  by  gentle  ignition.  Then  moisten  the  substance  with 
concentrated  hydrochloric  acid  and  dissolve  it  in  a  little  hot 

217 


218  COMPOUNDS   OF  THE  RARE   ELEMENTS. 

water;  again  evaporate  the  solution  to  dryness,  and  dehydrate 
the  residue  completely  by  heating  to  about  160°.  Extract  the 
residue  with  absolute  alcohol  either  in  a  Soxhlet  apparatus  or  in 
a  simple  flask  with  a  return  condenser.  In  the  latter  case,  extract 
eight  times  and  distil  each  of  the  alcoholic  solutions  poured  from 
the  insoluble  residue,  using  the  alcohol,  thus  recovered,  for  the 
next  extraction.  Finally,  after  all  of  the  alcohol  has  been  dis- 
tilled from  the  extract,  dissolve  the  salt  in  a  little  water  and 
precipitate  lithium  carbonate  by  adding  ammonium  carbonate 
and  a  little  ammonia.  From  lepidolite  1  to  3  g.  of  pure  lithium 
carbonate  should  be  obtained,  from  petalite  a  little  more,  and 
from  spodumene  about  twice  as  much.  Test  the  purity  of  the 
preparation  by  means  of  the  spectroscope. 

The  insoluble  salt  left  by  the  alcoholic  extraction  contains 
rubidium,  besides  other  of  the  alkali  metals.  Dissolve  it  in 
water,  treat  the  solution  with  about  1  c.c.  of  10%  chlorplatinic 
acid  solution,  and  allow  it  to  stand  over  night.  Next  morning 
filter  off  the  insoluble  chlorplatinates  and  ignite  them  in  a  Rose 
crucible  in  a  current  of  hydrogen.  Dissolve  the  residue  in  a 
little  water  and  test  it  with  the  spectroscope  for  rubidium. 

Spectroscopic  Analysis. 

The  scale  of  the  spectroscope,  which  is  usually  graduated  arbi- 
trarily, must  first  be  standardized  according  to  the  wave  lengths 
of  light.  After  the  apparatus  is  properly  set  up,  make  the 
sodium  line  fall  upon  a  certain  division  of  the  scale;  or,  in  case 
the  scale  telescope  is  not  movable,  observe  the  position  of  this 
line  as  accurately  as  possible.  Then  note  the  position  of  the 
following  lines: 


Wave  Length 

Wave  Length 

Kared 

768  X  10-6  mm. 

Srs  blue 

461  X  10-8mm. 

Lia  red 

671 

Ha  red 

656 

Na  yellow 

589 

H0  blue 

486 

Tl  green 

535 

Hv  violet 

434 

The  hydrogen  lines  are  very  sharp  and  easy  to  observe;  ready 
prepared  tubes  filled  with  hydrogen  under  diminished  pressure 
can  be  used,  and  are  to  be  excited  by  means  of  an  induction  coil. 
Plot  the  scale  readings  as  abscissas  and  the  corresponding  wave 


BERYLLIUM   HYDROXIDE.  219 

lengths  as  ordinates,  and  connect  the  separate  points  by  a  smooth 
curve.  One  millimeter  on  the  plot  may  be  taken  to  represent 
one  scale  division,  and  one-half  millimeter  on  the  other  axis  to 
represent  one  unit  of  wave  length.  From  this  curve  the  wave 
length  may  be  read  that  corresponds  to  each  division  on  the 
scale. 

If  the  lithium  preparation  is  pure  it  shows  only  the  line  671, 
and  eventually  the  weaker  line  at  610,  when  subjected  to  the 
spectroscopic  test.  If  it  still  contains  calcium  a  red  line  appears 
at  616,  and  green  bands  at  about  500.  Rubidium  is  recognized 
by  red  lines  at  795  and  780;  caesium  by  the  yellowish-red  lines 
622,  601,  584,  and  the  blue  lines  459  and  456. 

163.    Beryllium  Hydroxide  from  Beryl. 

Beryl  is  essentially  a  beryllium  aluminum  silicate,  3  BeO.Al2O3.6  SiO2, 
and  is  most  readily  decomposed  by  treatment  with  acid  ammonium  fluoride, 
whereby  silicon  fluoride  escapes  as  a  gas,  aluminum  goes  into  the  difficultly 
soluble  aluminum  fluoride,  and  beryllium  into  the  soluble  beryllium  fluoride. 
The  beryllium  fluoride  is  transformed  into  the  sulphate,  and  the  latter  is  dis- 
solved in  water  and  freed  from  traces  of  aluminum  by  means  of  concentrated 
ammonium  carbonate  solution.  Beryllium  remains  in  the  solution  as  a  double 
compound  with  ammonium  carbonate.  The  small  amount  of  iron  present  is 
precipitated  by  adding  ammonium  sulphide,  and  the  beryllium  is  finally 
thrown  from  the  filtrate  in  the  form  of  hydroxide. 

Treat  50  g.  of  finely  powdered  beryl  in  a  large  platinum  dish 
with  200  g.  of  neutral  ammonium  fluoride  and  300  g.  of  concen- 
trated hydrofluoric  acid,  and  heat  the  mixture  under  the  hood 
on  a  Babo  funnel  until  thick  vapors  escape  freely.  Continue 
the  heating  and  gradually  increase  the  temperature  until  the 
escape  of  vapors  practically  ceases,  allow  the  porous  mass  to 
cool,  rub  it  to  a  powder  with  a  platinum  spatula  or  a  wooden 
stick,  and  heat  again  until  nothing  more  is  vaporized.  Boil  the 
residue,  which  amounts  to  about  37  g.,  five  times  with  water, 
allowing  the  sediment  to  settle  and  decanting  the  clear  liquid 
through  a  plaited  filter  after  each  extraction.  About  15  g.  of 
aluminum  fluoride  remain  undissolved  after  this  treatment. 
Evaporate  the  filtrate  in  a  platinum  dish,  heat  the  residue  with 
30  g.  of  concentrated  sulphuric  acid  until  thick  fumes  of  sulphur 
trioxide  escape,  and  dissolve  the  dry  sulphate  in  100  to  200  c.c.  of 
water.  Stir  this  solution  slowly  into  a  cold  solution  of  150  g.  of 


220  COMPOUNDS  OF  THE  RARE  ELEMENTS. 

ammonium  carbonate  in  600  c.c.  of  water,  whereby  only  a  slight 
clouding  should  occur.  Then  dilute  with  an  equal  volume  of 
boiling  water  and  add  2  to  3  c.c.  of  ammonium  sulphide,  whereby 
any  iron  present  —  which  is  usually  but  little  —  gives  a  voluminous 
precipitate  of  iron  sulphide.  Filter  through  a  large  plaited  filter, 
the  point  of  which  is  reinforced  by  linen  cloth  folded  under 
it.  Wash  the  precipitate  with  hot  water  containing  a  little 
ammonium  sulphide,  and  concentrate  the  filtrate  to  one-half. 
The  beryllium  precipitates  as  beryllium  hydroxide  during  the 
evaporation.  Collect  the  precipitate  on  a  filter  and  boil  down 
the  solution  still  further,  thus  obtaining  the  little  beryllium 
which  is  left  as  a  less  pure  product  than  the  first.  Wash  the 
hydroxide  with  water  and  dry  it  at  a  gentle  heat.  Yield,  about 
13  g.  By  igniting,  5  to  6  g.  of  beryllium  oxide  could  be  obtained 
as  a  loose  white  powder,  but  for  use  in  the  following  preparation 
the  hydroxide  should  be  taken. 

If  a  large  platinum  dish  is  not  available  for  the  above  work,  a 
sheet-iron  crucible  can  be  used. 

164.    Basic  Beryllium  Acetate,  Be4O(CH3CO2)6. 

Beryllium  hydroxide  and  acetic  acid  form  a  stable  and  very  remarkable 
compound  of  the  above  composition.  It  melts  undecomposed  sharply  at 
289.5,  and  boils,  also  without  decomposition,  at  342  to  343°.  This  acetate 
can  be  used  to  purify  beryllium  preparations  from  all  other  metals,  and  it 
can  be  obtained  of  such  a  definite  composition  and  in  such  a  high  degree  of 
purity  that  it  has  been  used  for  determining  the  atomic  weight  of  beryllium.1 

Dissolve  the  beryllium  hydroxide  obtained  in  No.  163  by 
warming  it  with  dilute  acetic  acid.  Evaporate  the  solution  on 
the  water  bath  and  dissolve  the  gummy  residue  in  about  200  c.c. 
of  glacial  acetic  acid  (Hood).  Heat  the  mixture  (still  under  the 
hood)  and  filter  the  solution  at  the  boiling  temperature;  the  salt 
crystallizes  on  cooling  in  beautiful  colorless  needles  or  octahedra. 
Concentrate  the  mother-liquor  to  one-fourth,  in  order  to  obtain 
the  small  remainder  of  the  salt,  and  dry  the  entire  product  in 
the  steam  closet. 

Determine  the  melting-point.  Boil  a  sample  of  the  prepara- 
tion in  a  test-tube;  the  vapors  condense  on  the  cooler  upper 
walls  to  form  a  white  crust,  and  almost  no  residue  remains  in  the 

1  C.  L.  Parsons:  J.  Am.  Chem.  Soc.  27,  721  (1904). 


COLUMBIUM   AND  TANTALUM   COMPOUNDS.  221 

bottom  of  the  tube.  The  salt  is  insoluble  in  cold  water  but  is 
decomposed  by  hot  water.  It  is  slightly  soluble  in  alcohol,  less 
so  in  ether,  but  dissolves  readily  in  chloroform. 


165.    Columbium  and  Tantalum  Compounds  from  Columbite. 

Columbium  and  tantalum,  the  most  important  elements  in  columbite  and 
tantalite,  are  separated  from  one  another  by  crystallizing  their  potassium 
double  fluorides.  Potassium  tantalum  heptafluoride,  K2  [TaF7],  is  but  spar- 
ingly soluble,  and  crystallizes  from  a  dilute  solution,  whereas  potassium 
columbium  oxyfluoride,  K2  [CbOF5],  separates  only  from  a  concentrated  solu- 
tion. This  classic  method  was  originated  by  Marignac  in  1866,  and  is  still 
used  almost  exclusively;  in  this  way  the  first  perfect  separation  of  the  two 
elements  was  obtained. 

Concerning  the  history  of  these  elements,  it  is  of  interest  to  note  that 
Hatchett,  in  1801,  first  obtained  from  American  columbite  a  peculiar  acid 
which  ten  years  later  was  shown  by  the  researches  of  Wollaston  to  be  identi- 
cal with  a  similar  acid  obtained  by  Eckeberg  in  1802  from  Swedish  minerals 
and  named  by  him  tantalic  acid.  The  influence  of  Berzelius  led  to  the  adop- 
tion of  the  name  tantalum  for  the  element.  By  more  careful  researches  dat- 
ing from  1844,  H.  Rose  found  in  columbite  a  new  element,  columbium,  and 
of  this  he  was  the  first  to  obtain  pure  compounds.  He  believed  for  a  time 
that  a  third  element,  which  he  named  pelopium,  was  also  present.  Matters 
were  further  complicated  by  the  assumption  by  various  investigators  of  other 
elements  in  this  group,  and  it  remained  for  the  work  of  Marignac  to  show 
that  all  of  the  minerals  worked  with  contained  both  columbium  and  tantalum 
in  varying  proportions,  and  that  the  preparations  studied  up  to  that  time 
were,  for  the  most  part,  mixtures. 

Heat  100  g.  of  potassium  bisulphate  gently  in  a  platinum  dish 
until  it  has  reached  a  state  of  quiet  fusion;  allow  the  mass  to 
just  solidify,  and  distribute  25  g.  of  finely  powdered  columbite 
over  its  surface.  Heat  again  with  slowly  rising  temperature 
until  finally  a  clear  melt  is  obtained.  Allow  the  mass  to  cool,  or 
chill  it  by  placing  in  cold  water;  break  up  the  solid  mass  and 
allow  it  to  stand  in  water  over  night.  Next  morning  pour  off 
the  clear  solution,  and  boil  out  the  residue  several  times  with 
water  containing  a  little  hydrochloric  acid.  Finally  collect  the 
residue  on  a  plaited  filter,  wash  it  with  hot  water  containing 
hydrochloric  acid,  and  dry  it  in  the  hot  closet. 

To  purify  this  product  completely,  fuse  it  again  with  potassium 
bisulphate  exactly  as  before  and  wash  the  residue  very  thoroughly 
with  the  dilute  acid.  Yield,  25  g.  of  a  powder  consisting  of  the 


222  COMPOUNDS  OF  THE  RARE  ELEMENTS. 

oxides  of  tantalum  and  columbium,  and  containing  30  to  50%  of 
water;  the  powder  is  usually  entirely  free  from  iron. 

Determine  the  actual  amount  of  the  oxides  present  by  igniting 
0.6  g.  of  the  powder  in  a  platinum  crucible  to  constant  weight. 
For  each  gram  of  the  anhydrous  oxides,  2.5  g.  of  concentrated 
hydrofluoric  acid  and  0.5  g.  of  potassium  carbonate  are  to  be 
used.  The  hydrofluoric  acid  must  be  pure  and  above  all  free 
from  fluosilicic  acid.  Test  for  the  presence  of  the  latter  with 
potassium  salt  solution  in  a  platinum  crucible. 

First  dissolve  the  oxides  in  the  hydrofluoric  acid,  and  pour  the 
solution  cold  through  a  plaited  filter  placed  in  a  funnel  which 
has  been  coated  with  paraffin.  Allow  the  filtrate  to  run  into  a 
platinum  dish  containing  a  filtered  solution  of  the  required 
amount  of  potassium  carbonate.  Dissolve  the  resulting  precipi- 
tate by  boiling  and  adding  sufficient  water,  then  evaporate  the 
solution  until  crystals  begin  to  deposit.  On  cooling,  fine  needles 
of  potassium  tantalum  fluoride  are  obtained.  After  twelve  hours 
collect  the  crystals  on  a  plaited  filter  in  a  paraffined  funnel 
and  wash  them  with  cold  water.  On  further  evaporation  of  the 
filtrate  the  remainder  of  this  salt  (which  is  but  very  little)  can 
be  obtained. 

Finally,  concentrate  the  solution  to  a  much  smaller  volume 
until,  instead  of  the  fine  needle-like  crystals,  larger  thin  plates  of 
potassium  columbium  oxyfluoride  are  obtained.  Several  further 
crops  of  these  crystals  should  be  obtained  from  the  mother-liquor. 

Recrystallize  both  of  the  salts  from  water  containing  some 
hydrofluoric  acid.  Yield,  usually  about  1.5  g.  K2  [TaT7]  and  30  g. 
K2[CbOF5];  but  the  relative  quantities  of  the  two  salts  vary 
according  to  the  composition  of  the  original  material. 

166.   Molybdenum  Compounds  from  Molybdenite. 

Molybdenum  occurs  in  nature  to  some  extent  in  the  form  of  molybdates 
(wulfenite,  PbMoO4,  and  powellite,  CaMoO4),  but  the  most  important  ore  is 
the  sulphide  (molybdenite,  MoS2). 

(a)  Ammonium •  Molybdate,  5(NH^2Mo04.7MoOs.2H20 

Grind  50  g.  of  molybdenite  and  sift  it  through  fine  wire  gauze; 
grind  again  all  the  powder  that  will  not  pass  through  the  sieve, 
and  continue  the  process  until  nothing  remains  behind.  Roast 


MOLYBDENUM  COMPOUNDS  FROM  MOLYBDENITE.   223 

the  powder  in  a  small  evaporating  dish,  heating  strongly  with  a 
Fletcher  burner,  and  stirring  frequently  until  the  mass  has  been 
largely  converted  into  yellow  molybdenum  trioxide  and  the  sul- 
phur has  escaped  as  sulphur  dioxide.  Carried  out  in  this  manner 
the  roasting  consumes  from  four  to  six  hours;  if  it  is  carried  out 
in  a  porcelain  tube,  in  which  a  better  circulation  of  air  is  obtained, 
the  same  result  is  accomplished  in  from  thirty  minutes  to  an 
hour.  The  tube  should  be  about  25  cm.  long  and  3  cm.  in  diam- 
eter, and  should  be  placed  in  a  slightly  inclined  position  within 
an  asbestos  heating  chamber  (Fig.  4,  p.  3). 

Extract  the  roasted  product  with  2-normal  ammonia  solution, 
dry  the  dark-colored  residue,  roast  it,  and  extract  it  again  with 
ammonia;  finally  repeat  the  process  once  more,  when  nothing 
but  a  grayish  gangue  remains. 

To  the  entire  ammoniacal  extract  add  three  drops  of  ammo- 
nium sulphide  to  precipitate  traces  of  copper;  after  standing 
twelve  hours,  filter,  add  a  drop  of  bromine  to  the  filtrate,  and 
evaporate  until  crystallization  takes  place,  adding  at  the  last  a 
few  drops  of  concentrated  ammonia.  The  addition  of  a  little 
alcohol  aids  the  separation  of  crystals.  Dry  the  product  in  the 
air.  Yield,  35  to  40  g.  of  small  flake-like  crystals;  theoretical 
yield  from  pure  molybdenite,  47.5  g. 

A  dilute  solution  of  ammonium  molybdate  slightly  acidified 
with  hydrochloric  acid,  gives  a  dark-blue  color  when  treated  with 
one  drop  of  stannous  chloride;  upon  further  addition  of  this 
reagent  it  becomes  a  dirty  green.  Dependent  preparation:  Molyb- 
denum Blue,  No.  24. 

(b)   Oxides  of  Molybdenum. 

Molybdenum  Trioxide,  Mo03.  Ignite  some  ammonium  molyb- 
date at  first  gently  and  then  strongly  in  an  evaporating  dish. 
The  ignited  product  usually  contains  some  lower  oxide;  to 
change  this  completely  to  the  trioxide,  place  the  material  in  a 
combustion  tube  and  heat  in  a  slow  current  of  oxygen,  using 
a  row  burner  (Fig.  2)  and  covering  the  tube  with  an  asbestos 
mantle.  Do  not  heat  the  substance  sufficiently  to  volatilize 
the  molybdenum  trioxide. 

Molybdenum  Dioxide,  Mo02.  Place  2  g.  of  molybdenum  triox- 
ide in  a  weighed  glass  tube,  about  35  cm.  long,  which  has  been 


224       COMPOUNDS  OF  THE  RARE  ELEMENTS. 

drawn  out  a  little  at  one  end,  and  fill  the  tube  with  hydrogen 
which  has  been  washed  successively  with  caustic  soda  solution, 
silver  nitrate,  potassium  permanganate,  and  sulphuric  acid. 
Heat  in  the  atmosphere  of  hydrogen  for  15  minutes  to  moder- 
ate redness  so  that  the  temperature  surely  exceeds  470°.  When 
cold,  determine  the  loss  in  weight,  and  if  necessary  repeat  the 
heating  until  finally  the  change  in  weight  corresponds  to  the 
change  from  the  trioxide  to  the  dioxide.  In  this  way  a  reddish- 
brown  glistening  powder  is  obtained.1  Stronger  ignition  in 
hydrogen  causes  the  formation  of  metallic  molybdenum. 

(c)   Chlorides  of  Molybdenum. 

Molybdenum  Dioxydichloride,  Mo02Cl2.  Place  2  g.  of  molyb- 
denum dioxide  in  a  glass  tube  about  50  cm.  long,  and  pass  over  it 
a  current  of  chlorine  which  is  dry  and  entirely  free  from  air. 
After  the  air  is  completely  expelled  from  the  tube,  heat  gently  by 
means  of  a  Bunsen  burner  with  a  flame  spreader.  The  volatile 
molybdic  acid  chloride  sublimes  and  deposits  as  a  loose  mass  of 
pinkish-white  plates  in  the  colder  parts  of  the  tube.  After  the 
reaction  is  ended  replace  the  chlorine  with  dry  carbon  dioxide,  and 
when  cold,  cut  the  tube  at  a  point  between  the  residue  and  the 
sublimed  product.  Transfer  the  crystals  immediately  to  a  glass- 
stoppered  tube,  since  they  are  very  hygroscopic,  and  make  the 
stopper  air-tight  with  a  little  vaseline. 

Molybdenum  Pentachloride,  MoCl5.  First  prepare  the  neces- 
sary metallic  molybdenum  by  heating  15  g.  of  molybdenum  tri- 
oxide in  a  current  of  hydrogen  as  hot  as  possible  in  a  combustion 
furnace  until  no  more  water  vapor  is  evolved.  This  method  of 
reduction  requires  about  three  hours,  and  even  then  it  is  expedi- 
ent to  pulverize  the  product  and  to  heat  it  a  second  time  in  hydro- 
gen. Molybdenum  reduced  by  the  Goldschmidt  method  (see 
footnote  below)  may  also  be  pulverized  and  used. 

Meanwhile  construct,  under  a  hood  with  a  strong  draft,  the 
apparatus  represented  in  Fig.  26.  Make  as  indicated  two  con- 

1  Molybdenum  dioxide,  being  non-volatile,  is  especially  suited  for  the 
aluminothermic  production  of  the  fused  metal.  Place  a  mixture  of  80  g.  of 
molybdenum  dioxide  and  21  g.  of  aluminum  powder  in  a  clay  crucible 
embedded  in  sand,  and  start  the  reaction  by  means  of  some  fuse  powder  (cf. 
No.  2).  Yield,  70  to  80%. 


MOLYBDENUM  COMPOUNDS  FROM  MOLYBDENITE.   225 

strictions  in  the  combustion  tube  so  that  at  these  points  the 
inside  diameter  is  from  1.0  to  1.2  cm.,  while  the  right-hand 
arm  of  the  tube  is  32  cm.  and  the  left  arm  60  to  70  cm.  long. 


fHood 


Burner 


Fig.  26. 

Place  about  6  g.  of  powdered  molybdenum  in  the  shorter  arm, 
and  introduce  carbon  dioxide  and  hydrogen  into  the  apparatus 
until  all  the  air  has  been  replaced  in  the  wash  bottles  and  the 
connections.  Then  close  the  pinch  cock  of  the  carbon  dioxide 
tube,  and,  while  passing  a  slow  current  of  hydrogen,  heat  the 
molybdenum  as  strongly  as  possible  with  a  row  burner  for  one 
or  two  hours,  preventing  loss  of  heat  by  using  the  asbestos 
mantle.  Drive  out  any  condensed  water  through  the  open 
end  of  the  tube,  C,  by  fanning  with  a  Bunsen  flame.  Allow 
the  apparatus  to  cool  completely  while  the  hydrogen  is  still 
passing.  Meanwhile  start  the  action  in  the  chlorine  generator. 
Pass  the  chlorine  gas  through  a  wash  bottle  containing  water, 
and  then  through  two  others  containing  concentrated  sulphuric 
acid;  at  the  outset,  keep  the  rubber  connector  which  admits  to 
the  combustion  tube  closed  by  means  of  a  pinch  cock,  and  allow 
the  chlorine  to  escape  as  indicated  through  the  side  arm  into  the 
ventilating  flue.  When  the  air  has  been  completely  driven  out 
of  the  evolution  flask  and  from  all  the  connections  (which,  with  a 
fairly  rapid  current  of  chlorine,  requires  about  an  hour),  replace 
the  hydrogen  in  the  combustion  tube  with  carbon  dioxide,  and 
then  the  latter  with  chlorine.  The  reaction  begins  either  of 
itself  or  on  heating  very  gently  with  the  row  burner.  Streams  of 
a  dull-red  vapor  rise  and  condense  beyond  the  constriction  at  B. 
By  heating  the  molybdenum  gently  with  the  row  burner  and 


226  COMPOUNDS  OF  THE  RARE   ELEMENTS. 

occasionally  playing  a  Bunsen  flame  over  the  constriction  at  B, 
bring  the  product  into  the  section  of  the  tube  BC  where  it  pre- 
cipitates in  a  shower  of  very  minute  crystals.  Avoid  heating  too 
strongly.  At  the  end  of  the  operation  only  a  few  gray  flocks 
remain  behind  to  the  left  of  B.  During  the  process  the  end  of 
the  tube  at  C  is  connected  with  a  wide  glass  tube  leading  into 
the  flue. 

Allow  the  crystals  to  cool  in  a  current  of  carbon  dioxide,  close 
C  with  a  cork,  and,  with  a  blast  lamp,  fuse  the  tube  together 
at  B.  By  repeatedly  tapping  the  tube  loosen  the  crystals  and 
transfer  them  to  a  preparation  tube  which  is  made  ready  as  fol- 
lows: Close  one  end  of  a  fusible  glass  tube  35  cm.  long  and  of 
the  same  width  as  the  combustion  tube,  and  make  a  constriction 
20  cm.  above  the  closed  end.  Fill  the  tube  with  carbon  dioxide, 
and  place  the  open  end  over  the  narrowed  end  C  of  the  com- 
bustion tube.  As  soon  as  the  substance  is  transferred  seal  the 
constricted  part  of  the  preparation  tube.  Dissolve  the  residue 
adhering  within  the  combustion  tube  in  alcohol;  considerable 
heat  is  evolved,  and  an  emerald-green  solution  is  obtained  which, 
when  treated  with  ammonia,  gives  a  grayish-brown  flocculent 
precipitate. 

Molybdenum  Trichloride,  MoCl3.  In  an  apparatus  like  that 
used  for  the  preparation  of  the  pentachloride,  substitute  a  com- 
bustion tube  prepared  as  shown  in  Fig.  27.  The  lengths  of  the 


Fig.  27. 

three  sections  of  the  tube  are  32  cm.,  8  cm.,  and  60  to  75  cm. 
respectively,  and  the  inside  diameter  at  the  constrictions  A,  B, 
and  C  is  from  1.0  to  1.2  cm.  Proceed  at  first,  exactly  as  directed 
above,  to  prepare  molybdenum  pentachloride  from  6  g.  of  molyb- 
denum, and  allow  the  greater  part  of  this  product  to  sublime 
into  the  longest  section  of  the  tube,  while  a  small  amount  con- 
denses in  the  8  cm.  section.  Allow  the  tube  to  cool,  replace  the 
chlorine  by  carbon  dioxide  and  the  latter  by  hydrogen.  Then 
place  the  combustion  tube  in  a  somewhat  inclined  position  so 
that  D  is  higher  than  A.  Heat  the  tube  at  C,  while  a  fairly 
strong  current  of  hydrogen  is  passing,  until  the  pentachloride 


TUNGSTEN   COMPOUNDS  FROM  WOLFRAMITE.  227 

there  vaporizes,  and  then  continue  to  heat  so  that  5  to  10  cm.  of 
the  section  CD  remains  constantly  filled  with  red  vapors.  At 
first  there  is  no  evidence  of  chemical  change;  but  after  some  time 
hydrogen  chloride  can  be  recognized  in  the  escaping  gases,  and, 
on  continuously  vaporizing  the  pentachloride,  which  repeatedly 
condenses  and  flows  back,  a  copper-red  film  is  deposited,  and 
gradually  a  mass  of  the  same  or  often  a  darker  color  is  produced. 
Carry  the  heating  gradually  towards  D  until  finally  all  the  pen- 
tachloride is  converted  to  the  trichloride;  two  to  three  hours  are 
required  for  this.  Too  strong  heating  is  to  be  avoided,  for  the 
lower  the  temperature  at  which  the  transformation  takes  place 
the  better  the  preparation.  Finally,  replace  the  hydrogen  by  car- 
bon dioxide  and  distil  the  remainder  of  the  pentachloride  from 
the  8  cm.  section  so  that  its  vapor  passes  over  the  trichloride. 
Allow  none  of  the  pentachloride  to  remain  with  the  preparation, 
but  drive  out  all  the  excess  at  D  after  removing  the  escape  tube 
which  leads  to  the  flue.  Allow  the  preparation  to  cool,  cut  the 
tube  into  several  pieces,  and  remove  the  crystals  with  a  glass 
rod.  Yield,  4  to  6  g.  The  product  is  not  hygroscopic ;  it  consists 
usually  of  a  crystalline  mass  the  color  of  red  phosphorus  but 
sometimes  of  feathery  crystalline  aggregates. 

167.   Tungsten  Compounds  from  Wolframite. 

The  most  common  tungsten  mineral  is  wolframite,  which  is  essentially  a 
mixture  of  manganous  and  ferrous  tungstates. 

(a)  Ammonium  Tungstate,  5  (NH4)2W04.7  W03. 11  H20. 

Boil  200  g.  of  finely  powdered  wolframite  gently  for  about 
three  hours  in  a  300-c.c.  Erlenmeyer  flask  with  50  c.c.  of  concen- 
trated hydrochloric  acid  and  10  c.c.  of  concentrated  nitric  acid; 
replace  from  time  to  time  the  acid  as  it  evaporates.  Then  dilute 
the  mass  with  a  large  amount  of  water  and  decant  off  the  solu- 
tion, which  contains  chiefly  ferric  and  manganous  chlorides,  from 
the  grayish-yellow  easy-settling  residue.  Again  boil  the  residue 
gently  for  two  hours  with  the  same  mixture  of  acids  and  repeat 
the  washing  and  decantation. 

Dilute  the  decanted  liquids  to  a  volume  of  1.0  to  1.5  liters, 
and  if  after  several  hours  a  yellow  precipitate  of  tungstic  acid 
separates,  collect  it  on  a  filter  after  pouring  off  most  of  the  liquid, 
and  add  it  to  the  main  insoluble  residue. 


228  COMPOUNDS   OF  THE  RARE  ELEMENTS. 

Wash  the  insoluble  residue,  consisting  of  silicic  acid,  unat- 
tacked  wolframite,  and  yellow  tungstic  acid,  first  with  water  con- 
taining a  little  hydrochloric  acid,  and  then  with  pure  water; 
transfer  the  powder  to  a  beaker,  and  warm  it  on  the  water  bath 
with  enough  2-normal  ammonia  so  that  all  of  the  tungstic  acid 
passes  into  solution.  Then  filter  and  treat  the  residue  with 
acids,  etc.,  exactly  as  at  first,  in  order  to  make  certain  that  no 
large  amount  of  the  original  mineral  remains  unattacked. 
Finally,  unite  all  the  ammoniacal  extracts,  filter  the  solution 
again,  and  concentrate  it  until  crystals  of  ammonium  tungstate 
begin  to  form.  Then  add  10  c.c.  of  concentrated  ammonia  and 
allow  the  solution  to  cool.  Drain  the  colorless  crystalline  meal 
on  a  filter,  and  work  up  the  mother  liquor  until  only  a  very  little 
remains.  Dry  the  preparation  in  the  hot  closet.  Yield,  about  90  g. 

A  solution  of  ammonium  tungstate  gives  "a  light-yellow  pre- 
cipitate with  stannous  chloride.  The  precipitate  dissolves  partly 
or  wholly  in  concentrated  hydrochloric  acid,  and,  on  heating  the 
latter  solution,  a  precipitate  is  thrown  down  which  at  first  is 
dirty  blue  and  later  becomes  pure  dark  blue. 

(6)   Tungstic  Acid  and  Tungsten  Trioxide. 

Yellow  Dibasic  Tungstic  Acid,  H2WO^.  Boil  3  g.  of  ammo- 
nium tungstate  gently  for  five  minutes  in  an  evaporating  dish 
with  a  mixture  of  10  c.c.  of  concentrated  hydrochloric  acid 
and  5  c.c.  of  concentrated  nitric  acid.  Allow  the  liquid  to  cool 
and  dilute  it  with  water.  Collect  the  solid  residue  on  a  suction 
filter,  wash  it  with  hot  water  until  the  washings  show  a  neutral 
reaction  to  blue  litmus,  and  dry  it  in  the  hot  closet.  By  igniting 
the  tungstic  acid  it  loses  one  molecule  of  water  and  is  transformed 
into  tungsten  trioxide.  Determine  the  loss  in  weight. 

White  Tungstic  Acid,  H4W05(?).  Dissolve  3  g.  of  ammonium 
tungstate  in  30  c.c.  of  water  with  the  addition  of  a  few  drops  of 
ammonia.  Then  add  to  the  solution  at  the  room  temperature  an 
equal  volume  of  10%  hydrochloric  acid.  White  tungstic  acid  is 
thereby  precipitated.  After  it  settles,  collect  the  solid  and  wash  it 
with  hot  water  containing  some  hydrochloric  acid.  If  the  filtrate 
is  opalescent  by  reflected  light,  it  contains  colloidal  tungstic  acid 
which  can  be  precipitated  by  gentle  heating.  Dry  the  prepara- 
tion at  a  moderate  temperature  on  top  of  the  hot  closet. 


WORKING  UP   OF   PITCHBLENDE.  229 

Tungsten  Trioxide,  W03.  This  oxide  can  be  obtained  by  heat- 
ing either  of  the  above  acids,  or  it  can  be  formed  directly  from 
ammonium  tungstate  by  heating  the  salt  in  a  porcelain  crucible 
at  first  over  a  Bunsen  flame  and  finally  over  a  blast  lamp.  The 
oxide  is  a  lemon-yellow  powder  which  in  the  sunlight  acquires  a 
greenish  tinge. 

(c)   Tungsten  Hexachloride,  WCIQ. 

First  prepare  metallic  tungsten,  as  a  grayish-black  powder,  by 
reducing  the  trioxide  in  an  atmosphere  of  hydrogen  at  the  highest 
temperature  obtainable  in  the  combustion  furnace.  (Cf .  the  prep- 
aration of  molybdenum  under  Molybdenum  Pentachloride,  p.  224.) 

Prepare  tungsten  hexachloride  from  the  metallic  tungsten 
according  to  the  directions  for  obtaining  molybdenum  penta- 
chloride  from  molybdenum  (p.  224).  Here  also  it  is  important 
to  exclude  moisture  and  atmospheric  oxygen  with  the  greatest 
care.  On  passing  chlorine  over  the  metal  some  yellow  tungsten 
oxychloride  is  formed  at  first;  expel  this  by  heating  the  tube  with 
a  free  flame,  and  allow  the  vapors  to  condense  only  when  they 
have  become  dark  colored.  The  hexachloride  is  a  dark-violet, 
very  hygroscopic,  crystalline  powder,  which  must  be  preserved 
in  a  sealed  tube.  Light  and  air  change  it  to  yellow  tungstyl 
chloride,  WO2C12. 


168.   Working  Up  of  Pitchblende  and  Testing  of  the  Components 

for  Radioactivity. 

Becquerel  discovered  in  1896  that  pitchblende  and  uranium  preparations 
send  out  peculiar  radiations  which,  like  the  Roentgen  rays,  can  be  detected 
by  their  action  upon  photographic  plates,  and  by  their  ability  to  make  the 
air  through  which  they  pass  a  conductor  of  electricity.  These  radiations 
are  called  Becquerel  rays,  and  the  substances  emitting  them  are  said  to  be 
radioactive.  By  separating  pitchblende  into  its  constituents  and  by  testing 
each  for  its  radioactivity,  M.  and  Mme.  Curie  (1898)  established  the  fact 
that  the  ability  to  emit  Becquerel  rays  becomes  concentrated  in  certain 
definite  constituents;  the  barium  sulphate,  which  they  prepared  from  the 
barium  residues  obtained  in  the  technical  working  up  of  the  ore,  was  found 
to  be  particularly  active;  and  from  this  they  succeeded  in  isolating  a  new 
element,  radium,  which  showed  to  the  very  highest  degree  the  property  of 
radioactivity.  The  lead,  bismuth,  uranium,  and  rare  earths  obtained  from 
pitchblende  are  also  radioactive,  although  to  a  lesser  degree. 


230  COMPOUNDS   OF   THE  RARE   ELEMENTS. 

Uranium  compounds  are  obtained  technically  from  pitchblende  by  digest- 
ing the  ore  with  nitric  acid,  evaporating  to  dryness,  and  extracting  the 
residue  with  water;  uranyl  nitrate  is  crystallized  from  the  solution.  The 
insoluble  material  is  then  extracted  with  a  large  amount  of  sodium  carbonate 
solution,  whereby  the  rest  of  the  uranium  is  dissolved  as  sodium  uranyl  car- 
bonate. The  residues  from  this  treatment  form  the  starting  material  for 
obtaining  radium. 

This  discovery  of  radium  and  radioactivity  was  epoch-making  for  both 
physics  and  chemistry.  So  intense  has  the  investigation  of  radioactivity 
become  that  special  journals  giving  original  literature  and  references  to  all 
work  performed  in  this  field  are  now  published  both  in  English  and  German. 

Pulverize  10  to  15  g.  of  pitchblende  as  finely  as  possible,  and 
digest  the  powder  for  several  hours  on  the  water  bath  with  a 
mixture  of  20  g.  of  concentrated  nitric  acid  and  an  equal  amount 
of  water.  Evaporate  the  mixture  to  dryness,  and  treat  the  resi- 
due twice  successively  with  5  g.  of  the  same  acid  mixture,  evapo- 
rating each  time  to  dryness.  Extract  the  dry  mass  with  water 
and  evaporate  the  filtrate  on  the  water  bath.  Place  the  dry 
residue  from  the  aqueous  solution  in  a  flask,  and  extract  it  several 
times  with  warm  ether  until  nothing  more  dissolves;  evaporate 
the  ethereal  solution  of  uranyl  nitrate,  taking  care  that  the 
vapor  does  not  take  fire.  Dissolve  the  salt  in  water,  precipitate 
it  with  ammonia,  and  ignite  the  precipitate  to  uranium  octo- 
oxide,  U308. 

Treat  with  aqua  regia  the  water-insoluble  residue  from  the 
nitric  acid  treatment,  and  evaporate  the  mass  to  dryness.  Moisten 
the  residue  with  a  little  concentrated  hydrochloric  acid,  extract 
it  with  hot  water,  and  combine  the  solution  with  an  aqueous  solu- 
tion of  the  residue  left  from  the  extraction  with  ether. 

Separate  the  constituents  of  the  solution  thus  obtained  accord- 
ing to  the  usual  procedure  of  qualitative  analysis.  The  working 
up  of  the  hydrogen  sulphide  precipitate  yields:  (1)  As,  Sb,  Sn; 
(2)  Pb;  (3)  Bi;  (4)  Cu.  Dissolve  the  ammonium  sulphide  pre- 
cipitate in  hydrochloric  acid,  oxidize  the  solution,  and  precipitate 
Fe,  Al,  and  the  rare  earths  with  ammonia.  Treat  this  pre- 
cipitate immediately  with  concentrated  sodium  carbonate  solu- 
tion, whereby  the  -remainder  of  the  uranium  dissolves ;  it  can  be 
precipitated  by  caustic  soda  from  this  solution  as  sodium  pyro- 
uranate.  Dissolve  the  residue  from  the  sodium  carbonate  treat- 
ment in  as  little  hydrochloric  acid  as  possible,  nearly  neutralize 


TESTING  FOR   RADIOACTIVITY.  231 

the  solution  with  ammonia,  and  then  treat  it  with  a  solution  of 
oxalic  acid.  This  precipitates  as  oxalate  any  thorium  that  is 
present.  On  treating  the  nitrate  with  ammonia,  the  hydroxides 
of  iron  and  aluminum  are  precipitated.  The  filtrate  from  the 
first  precipitation  with  ammonia  may  contain  zinc  and  cobalt; 
these  metals  should  be  precipitated  with  ammonium  sulphide, 
but  they  need  not  be  separated  from  each  other. 

The  filtrate  from  the  ammonium  sulphide  group,  on  being 
treated  with  ammonium  oxalate,  yields  a  precipitate  of  oxalates 
of  the  alkaline  earth  metals. 

That  part  of  the  ore  which  does  not  dissolve  in  nitric  acid  nor 
during  the  subsequent  treatment  with  aqua  regia,  should  be  fused 
in  a  porcelain  crucible  with  a  mixture  of  sodium  and  potassium 
carbonates.  Extract  the  fusion  with  water,  and  wash  the  residue 
until  the  filtrate  no  longer  gives  a  test  for  sulphate.  Dissolve 
the  carbonates  in  dilute  nitric  acid,  and  test  the  solution  for  lead 
and  the  alkaline  earth  metals.  The  combined  alkaline  earths 
should  be  separated  from  each  other  only  when  larger  amounts  of 
pitchblende  are  worked  up. 

Testing  for  Radioactivity. 

(a)  Photographically.  Fold  a  piece  of  black  paper  around  a 
photographic  dry  plate  (in  the  dark  room),  so  that  the  sensitive 
side  is  covered  with  one  thickness  of  the  paper,  and  place  the 
whole  in  a  box  with  the  sensitive  side  up.  Upon  the  plate 
arrange  samples  of  the  original  pitchblende  and  of  the  different 
preparations  obtained  from  it,  each  enveloped  in  a  piece  of  paper. 
Record  the  position  of  each  specimen  and  then  close  the  box. 
Open  it  at  the  end  of  twenty-four  hours  and  develop  the  plate  in 
the  usual  manner.1 

(fr)  Electroscopically .  An  approximate  measure  of  activity 
can  be  obtained  with  the  aid  of  a  sensitive  gold-leaf  electroscope. 
A  scale  should  be  placed  so  that  the  distance  between  the  gold 
leaves  can  be  read.  The  top  of  the  electroscope  should  consist  of 
a  metal  plate,  and  above  this  at  a  definite  distance  a  second  metal 

1  Uranyl  nitrate,  when  prepared  as  above  by  the  ether  method,  does  not 
show  its  full  activity  at  once;  its  maximum  intensity  is  "recovered"  very 
slowly,  —  one-half  in  twenty-two  days. 


232  COMPOUNDS  OF   THE   RARE  ELEMENTS. 

plate  that  is  connected  with  the  earth  should  be  supported. 
Charge  the  electroscope,  and  measure  the  conductivity  of  the 
layer  of  air  between  the  two  plates  by  determining  the  time 
required  for  the  two  leaves  of  the  electroscope  to  reach  the  zero 
position.  It  is  sufficient  to  observe  over  how  many  scale  divi- 
sions the  leaves  pass  within  a  certain  length  of  time,  e.g.,  the 
scale  divisions  per  minute.  To  measure  the  activity,  spread  an 
amount  of  the  substance,  weighed  to  an  accuracy  of  two  figures, 
upon  a  piece  of  paper,  and  place  the  paper  on  the  lower  metal 
plate.  Adjust  the  second  plate  in  position,  charge  the  electro- 
scope, and  make  the  readings.  The  quotient  obtained  by  divid- 
ing the  number  of  scale  divisions  traversed  per  minute  by  the 
weight  of  the  substance,  gives  the  specific  activity.  The  observed 
rate  may  be  corrected  by-  subtracting  from  it  the  velocity  with 
which  the  leaves  come  together  in  a  blank  experiment  with  none 
of  the  active  preparation.  If  it  is  desired  to  compare  the  values 
obtained  with  those  that  have  been  published  in  the  literature, 
the  apparatus  may  be  standardized  by  using  uranyl  nitrate. 

Of  the  active  radiations,  the  part  consisting  of  the  so-called 
a-rays  is  held  back  by  a  sheet  of  paper.  If,  therefore,  the  experi- 
ment is  repeated  exactly  as  above,  except  that  the  preparation  is 
covered  with  filter  paper,  the  value  then  obtained  corresponds  to 
the  activity  of  the  rays  that  pass  through  the  paper,  that  is,  of 
the  so-called  /?-rays.  The  difference  between  the  two  values  gives 
the  activity  of  the  a-rays.  Radioactive  lead  preparations  emit 
principally  a-rays. 

169.   Uranium  Compounds. 

As  the  starting  material  in  the  preparation  of  uranium  compounds,  either 
uranium  nitrate  prepared  from  pitchblende  according  to  the  following  pro- 
cedure, or  the  commercial  product,  may  be  used. 

Uranyl  Nitrate,  U02(N03}2.  Heat  50  g.  of  finely  powdered 
and  sifted  pitchblende  in  an  evaporating  dish  with  150  c.c.  of 
30%  nitric  acid  until  the  excess  of  acid  has  been  expelled. 
Extract  the  mass  with  hot  water,  and  evaporate  the  filtered  solu- 
tion to  dryness  on  the  water  bath.  Boil  the  dry  residue  repeat- 
edly with  ether  in  a  small  flask,  and  evaporate  the  ethereal  extract 
to  dryness,  taking  great  care  that  the  vapor  does  not  take  fire.1 
Recrystallize  the  last  residue  from  water.  Yield,  about  35  g. 

1  Cf.  footnote  2,  p  214. 


URANIUM  COMPOUNDS.  233 

Uranyl  Hydroxide,  U02(OH)2. 

The  hydroxide  of  hexavalent  uranium  is  amphoteric,  that  is,  it  has 
acidic  as  well  as  basic  properties.  With  acids  it  forms  the  uranyl  salts,  e.g., 
UO2(NO3)2;  with  bases  it  gives  the  difficultly  soluble  salts  of  pyro-uranic 
acid,  e.g.,  K2U2O7.  For  this  reason  the  free  uranyl  hydroxide  cannot  be 
obtained  by  adding  a  solution  of  an  alkali  hydroxide  to  one  containing  a 
uranyl  salt;  it  is  prepared  by  heating  uranyl  nitrate  with  anhydrous  alcohol. 

Heat  20  g.  of  uranyl  nitrate  with  50  g.  of  absolute  alcohol  on  the 
water  bath  in  such  a  way  that  the  alcohol  evaporates  slowly, 
but  does  not  boil.  After  some  time  yellow  uranyl  hydroxide 
separates;  add  more  alcohol  and  evaporate  further.  Finally, 
extract  the  precipitate  with  water,  and  dry  it  at  a  moderate 
temperature. 

Alkali  Pyro-uranates.  To  a  solution  of  uranyl  nitrate,  add 
potassium  hydroxide,  sodium  hydroxide,  or  ammonia  until  all  of 
the  uranium  is  just  precipitated.  Drain  the  precipitate,  wash  it 
carefully  with  water  and  dry  it  in  the  hot  closet.  Sodium  pyro- 
uranate  is  prepared  industrially,  and  is  used  in  the  manufacture 
of  yellow-green  fluorescent  glass. 

Ammonium  Uranylcarbonate,  U02C03  •  2  (NH4)2C03.  Add  a 
solution  of  ammonium  carbonate  carefully  to  a  dilute  solution 
containing  20  g.  of  uranyl  nitrate  until  precipitation  is  just  com- 
plete; an  excess  of  the  reagent  dissolves  the  precipitate.  Filter 
off  the  small,  light-yellow  crystals,  and  wash  them  successively 
with  water,  alcohol,  and  ether.  On  standing,  the  preparation 
loses  ammonium  carbonate. 

Uranium  Trioxide,  U0y  Dry  some  uranyl  ammonium  carbon- 
ate, or  uranyl  hydroxide,  at  100°;  then  place  it  in  a  test-tube,  and, 
while  shaking,  heat  it  in  an  oil  or  paraffin  bath  at  250°  to  300° 
until  the  color  has  become  brick-red. 

Uranium  Octo-oxide,  U308.  This  oxide  is  obtained  by  ignit- 
ing uranium  trioxide  (or  any  oxide  of  uranium),  or,  more  con- 
veniently, ammonium  pyro-uranate,  in  a  porcelain  crucible;  it  is 
a  dark-green  powder. 

Uranous  Oxalate,  C/(C204)2. 

Uranyl  salts  are  changed  by  reduction  into  uranous  salts.  The  latter, 
however,  are  easily  oxidized  again,  although  the  difficultly  soluble  uranous 
oxalate  is  relatively  stable.  A  satisfactory  reducing  agent  is  the  sodium  salt 
of  hyposulphurous  acid,  which  can  be  procured  in  the  form  of  a  paste  contain- 
ing 50%  of  the  salt. 


234  COMPOUNDS  OF   THE   RARE  ELEMENTS. 

Prepare  a  solution  containing  25  g.  of  the  reducing  agent,  and 
add  this  to  a  solution  of  25  g.  of  uranyl  nitrate  in  100  c.c.  of -water 
until  the  solution  is  decolorized;  a  precipitate  is  formed  which  is 
at  first  brown  in  color  but  changes  to  a  lighter  shade.  Dissolve 
the  precipitate  in  hydrochloric  acid,  filter,  and  treat  the  filtrate 
with  a  hot  solution  of  15  g.  of  oxalic  acid  in  120  c.c.  of  water. 
About  26  g.  of  gray,  finely-crystalline  uranous  oxalate  are  pre- 
cipitated. 

To  purify  the  precipitate,  dissolve  it  in  a  solution  of  neutral 
ammonium  oxalate,  using  6  g.  of  the  latter  salt  and  100  c.c.  of 
water  for  each  10  g.  of  uranous  oxalate.  Reprecipitate  the  salt 
from  the  filtered  solution  by  adding  hydrochloric  acid.  After 
some  hours  filter  off  the  precipitate,  wash  it  with  water  and  with 
alcohol,  and  dry  it  in  the  hot  closet.  Yield,  9  g.  from  each  10  g. 
of  the  impure  salt. 

Ammonium  Urano-oxalate,  (NH^)4[U(C204)4].  Digest  10  g.  of 
uranous  oxalate  with  a  solution  of  5  g.  of  neutral  ammonium  oxa- 
late in  50  c.c.  of  water;  filter  off  the  excess  of  uranous  oxalate, 
and  precipitate  the  double  salt  from  the  filtrate  by  adding  alco- 
hol a  little  at  a  time.  Allow  the  mixture  to  stand  for  one  or  two 
days,  and  then  filter  off  the  crystal  meal  which  has  by  that  time 
become  coarser.  Yield,  12  g. 

Uranium  Tetrachloride,  UC14.  Uranium  tetrachloride  is  pre- 
pared in  the  apparatus  described  under  No.  166  (c),  by  passing  per- 
fectly dry  chlorine  over  a  mixture  of  uranium  oxide,  U30S,  and 
one-eighth  of  its  weight  of  ignited  wood  charcoal,  the  whole 
being  heated  as  hot  as  possible  in  a  combustion  furnace.  The 
principle  is  the  same  as  that  outlined  under  No.  52.  Uranium 
tetrachloride  condenses  in  the  front  part  of  the  tube,  and  is  freed 
from  admixed  pentachloride  by  heating  it  in  a  current  of  carbon 
dioxide  at  about  150°.  In  this  way  a  very  hygroscopic  mass  of 
greenish-black  crystals  is  obtained.  The  tetrachloride  dissolves 
in  water  with  evolution  of  heat,  and  yields  a  green  solution  (cf. 
Uranous  Oxalate). 

170.   Thorium  Compounds  from  Monazlte. 

Thorium  was  first  discovered  by  Berzelius  in  the  rare  Scandinavian  min- 
erals thorite  and  orangite.  When  the  oxide  of  thorium  became  of  industrial 
importance  through  its  use  in  the  Welsbach  incandescent  mantles,  a  more 


THORIUM  COMPOUNDS   FROM   MONAZITE.  235 

abundant  source  was  sought.1  This  was  found  in  the  mineral  monazite,  a 
phosphate  of  the  rare  earths  and  of  thorium,  which  contains  4  to  7%  of  thorium 
oxide  and  50  to  60%  of  the  rare  earths,  about  one-half  of  the  latter  being 
cerium  oxide.  This  mineral  occurs  frequently,  although  only  in  small 
amounts,  in  primary  rocks;  but  it  is  found  in  some  places  concentrated  in 
secondary  deposits. 

Add  100  g.  of  finely  powdered  monazite  sand,  a  little  at  a  time, 
to  150  g.  of  concentrated  sulphuric  acid  which  is  heated  to  200° 
in  an  evaporating  dish  upon  a  Babo  funnel.  Keep  the  mass  at 
this  temperature  for  half  an  hour  after  all  of  the  mineral  is  added; 
then  allow  it  to  cool  completely,  and  pour  it  very  slowly  with 
constant  stirring  into  300  c.c.  of  water,  whereby  the  temperature 
must  not  be  allowed  to  rise  above  25°  at  any  time.  Filter  off  the 
residue,  dry  it,  and  repeat  the  above  treatment,  using  50  to  75  g. 
of  sulphuric  acid,  and  subsequently  adding  the  mixture  to  150  c.c. 
of  water. 

Separation  of  Thorium.  Unite  the  two  solutions  without  dilut- 
ing them  unnecessarily,  and  add  a  solution  of  50  g.  of  oxalic 
acid  in  500  c.c.  of  water  until  no  further  precipitate  forms. 
Filter  off  the  precipitate,  and  wash  it  first  with  water  containing 
small  amounts  of  oxalic  and  sulphuric  acids,  and  finally  with  a 
little  pure  water.  Yield,  about  75  g. 

Make  the  filtrate  approximately  neutral  with  sodium  carbon- 
ate and  again  add  oxalic  acid  to  throw  out  thorium.  This  second 
precipitate  is  of  one-third  to  one-fourth  the  quantity  of  the  first, 
and  is  not  entirely  free  from  phosphates.  It  is  well  to  set  it 
aside  and  to  work  it  up  together  with  the  fresh  mineral  when 
making  the  next  preparation. 

Boil  up  the  moist  oxalates  with  a  solution  of  230  g.  of  anhy- 
drous sodium  carbonate  in  one  liter  of  water;  the  thorium  dis- 
solves as  sodium  thorium  carbonate,  while  the  rare  earths  remain 
behind  as  carbonates.  Filter  the  warm  liquid  immediately,  and 
acidify  the  filtrate  with  hydrochloric  acid;  3  to  4  g.  of  thorium 
oxalate  are  precipitated. 

In  order  to  recover  the  remainder  of  the  thorium  from  the 
mixture  of  insoluble  carbonates,  dissolve  the  latter  in  just  the 
necessary  amount  of  nitric  acid,  evaporate  the  solution  to  dry- 

1  It  is  interesting  to  note  that  although  one  kilo  of  thorium  nitrate  was 
worth  about  $500  in  1894,  the  price  had  fallen  to  $7  in  1900. 


236  COMPOUNDS  OF  THE  RARE   ELEMENTS. 

ness  on  the  water  bath,  dissolve  the  residue  in  hot  water,  and 
precipitate  thorium  hydroxide  from  the  boiling  solution  by  add- 
ing sodium  thiosulphate.1 

To  precipitate  the  rare  earths,  add  sodium  carbonate  to  the 
filtrate  from  the  thorium  hydroxide,  drain  and  wash  the  pre- 
cipitate, and  dry  it  in  the  hot  closet.  This  product  may  be 
worked  up  according  to  No.  171. 

Purification  of  the  Thorium.  Dissolve  the  thorium  precipi- 
tates in  hydrochloric  acid,  adding,  if  necessary,  a  little  nitric 
acid;  evaporate  the  solution  to  dryness,  take  up  the  residue  in 
water,  filter  if  necessary,  and  treat  the  solution  at  60°  with  a 
solution  of  sodium  thiosulphate.  Collect  the  precipitate  on  a 
filter,  wash  it  with  water  containing  some  ammonium  nitrate  (as 
it  has  a  tendency  to  pass  into  colloidal  solution),  and  ignite  it  to 
thorium  dioxide. 

Thorium  Sulphate.  Heat  the  thorium  oxide  with  concen- 
trated sulphuric  acid  in  a  porcelain  crucible  until  the  excess  of 
acid  has  been  expelled,  moisten  the  residue  with  a  few  drops 
more  of  sulphuric  acid  and  heat  it  as  before,  but  avoid  bringing 
it  to  a  red  heat.  Pulverize  the  thorium  sulphate,  which  should 
be  free  from  acid  salt,  and  add  it  gradually,  with  vigorous 
stirring,  to  five  times  its  weight  of  ice-water.  If,  after  a  little 
while,  a  considerable  residue  remains  undissolved,  remove  it  and 
subject  it  again  to  the  treatment  with  concentrated  sulphuric 
acid. 

By  warming  the  filtered  solution  to  30  —  35°,  thorium  sulphate 
octohydrate,  Th(SO4)2  •  8  H2O,  together  with  a  little  enneahydrate, 
Th(SO4)2  •  9  H2O,  is  caused  to  separate.  Maintain  the  solution 
at  this  temperature,  and  allow  it  to  evaporate  until  all  of  the 
salt  has  separated.  Drain  the  crystals,  wash  them  with  a 


1  The  salts  of  tetravalent  thorium  are  more  easily  hydrolyzed  than  those 
of  the  trivalent  rare-earth  metals.  Thus  thorium  hydroxide  is  precipitated 
while  the  salts  of  the  trivalent  metals  remain  unchanged.  The  principle  of 
this  separation  is  similar  to  that  of  the  basic  acetate  method  used  in  analyti- 
cal chemistry. 

According  to  another  method,  the  solution  is  warmed  to  60  to  70°  together 
with  an  excess  of  a  10%  hydrogen  peroxide  solution;  the  thorium  and  some 
cerium  are  thrown  down  in  the  form  of  a  flocculent  precipitate  which  can  be 
easily  filtered. 


SEPARATION  OF  THE  RARE   EARTHS.  237 

little    water,    then   with   alcohol,    and  dry    them    at    the    room 
temperature. 

Atomic  Weight  Determination  of  Thorium.  In  order  to  illus- 
trate the  principle  of  an  atomic  weight  determination  by  the  sul- 
phate method,  prepare  first  some  anhydrous  thorium  sulphate 
from  about  2  g.  of  the  above  hydrated  salt.  Heat  the  crystallized 
salt  to  400°  in  a  weighed  platinum  crucible  which  is  supported  by 
a  platinum  triangle  within  a  larger  crucible.  When  the  weight 
has  become  constant,  the  amount  of  the  anhydrous  thorium  sul- 
phate, a,  is  given  by  subtraction.  Then  ignite  the  crucible  over  the 
blast  lamp  until  the  weight  has  again  become  constant.  This  gives 
the  weight  of  thorium  dioxide,  b.  If  the  atomic  weight  of  oxygen 
is  taken  at  16.00  and  that  of  sulphur  as  32.06,  the  atomic  weight 
of  thorium  is  obtained  by  solving  for  x  in  the  expression: 

x  +2  •  32.06  +  8*  16.00  _  a 
x  +  2  •  16.00  ~b' 


171.    Separation  of  the  Rare  Earths. 

The  rare  earths  form  a  group  of  very  closely  related  sesquioxides,  the 
separation  and  characterization  of  which  for  a  long  time  offered  considerable 
difficulties.  The  properties  of  the  analogous  compounds  of  these  earths  do 
not  differ  sharply  enough  from  each  other  to  permit  a  complete  separation 
to  be  made  in  a  single  operation;  but  the  slight  gradations  in  the  properties 
may  in  general  be  used  to  effect  a  satisfactory  separation  if  a  given  pro- 
cess is  systematically  repeated  again  and  again.  When  these  slight  grada- 
tions are  taken  into  account,  the  rare  earths  fall  into  a  classification  which 
corresponds  closely  to  their  occurrence  in  nature. 

The  first  distinction  is  made  between  the  cerium  earths  and  the  yttrium 
earths.  The  latter  were  discovered  by  Gadolin  at  the  end  of  the  eighteenth 
century  in  a  mineral  found  in  a  feldspar  quarry  near  Ytterby,  in  Sweden,  and 
afterwards  named  after  him  gadolinite.  The  cerium  earths  were  discovered 
at  the  beginning  of  the  last  century  by  Berzelius  and  Hisinger  in  the  mineral 
cerite  which  had  previously  been  investigated  by  Scheele  without  success. 
Soon  afterwards  lanthanum  and  didymium  were  discovered  in  the  first  half  of 
the  nineteenth  century  by  Mosander,  a  pupil  and  friend  of  Berzelius.  This 
same  investigator  found  that  yttrium  is  accompanied  by  terbium  and  erbium. 
The  separation  of  the  earths  was  accomplished  by  two  different  processes 
which  in  a  more  perfected  form  serve  even  to-day  as  the  standard  methods. 
Of  these  processes,  one  rests  upon  differences  in  the  basicity  of  the  earths,  as 
manifested  in  the  varying  hydrolytic  and  thermic  dissociation  of  their  salts ; 
the  second  process  rests  upon  differences  in  the  solubility  of  the  double  salts. 


238 


COMPOUNDS  OF   THE  RARE  ELEMENTS. 


The  cerium  earths  are  more  weakly  basic  and  form  more  difficultly  soluble 
potassium  double  sulphates  than  the  yttrium  earths.  Cerium  itself  occu- 
pies a  characteristic  position,  as  it  is  the  only  member  of  the  group  from 
which  two  series  of  salts  are  derived:  the  cerous  salts  with  trivalent  cerium, 
and  the  eerie  salts  containing  tetravalent  cerium. 

For  the  qualitative  characterization  of  the  earths,  the  color  of  their  salts 
is  first  of  all  of  importance.  Lanthanum  and  cerous  salts  are  colorless,  eerie 
salts  are  yellowish-red,  erbium  salts  pink,  didymium  salts  violet.  It  marked 
then  the  beginning  of  a  new  era  in  the  history  of  the  rare  earths,  when,  in 
1861,  through  the  application  of  spectrum  analysis,  it  became  possible  to 
measure  exactly  the  color  of  salt  solutions  by  means  of  their  absorption  spec- 
tra, and  the  color  of  glowing  vapors  by  means  of  emission  spectra.  It  was  in 
this  epoch  that  the  discovery  of  the  periodic  system  of  the  elements  was 
made  (1869),  by  the  aid  of  which  Mendelejeff  first  recognized  the  trivalency 
of  the  rare-earth  metals  and  predicted  the  existence  of  the  element  scandium, 
which  was  later  discovered  by  Nilson.  Further  investigation  and  the  exami- 
nation of  new  minerals  has  since  that  time  added  a  number  of  elements  to 
the  group.  In  the  presence  of  didymium  Lecoq  de  Boisbaudran  discovered 
samarium.  In  the  group  of  the  yttrium  earths,  which  since  the  time  of 
Mosander  was  studied  especially  by  Bahr  and  Bunsen,  ytterbium  was  dis- 
covered by  Marignac,  while  Cleve  added  the  elements  holmium  and  thulium. 
Marignac  also  discovered  gadolinium,  which  together  with  terbium  and  euro- 
pium occupies  an  intermediate  position  between  the  cerium  and  yttrium 
earths.  Finally,  didymium  was  separated  by  von  Welsbach  in  1885  into 
praseodymium,  the  salts  of  which  are  green,  and  neodymium,  whose  salts  are 
violet. 

The  most  recent  epoch  dates  from  the  technical  application  by  von  Wels- 
bach of  thorium  and  cerium  oxides  in  the  incandescent  gas-lighting  industry. 

The  following  table  gives  a  summary  of  the  rare  earths  and  their  atomic 
weights: 


Cerium  earths. 

Terbium  earths. 

Yttrium  earths. 

Lanthanum,  La           139.0 
Cerium,  Ce                   140.25 
Praseodymium,  Pr      140.6 
Neodymium,  Nd          144.3 
Samarium,  Sm             150.4 

Europium,  Eu        152.0 
Gadolinium,  Gd     157.3 
Terbium,  Tb          159.2 

Scandium,  Sc        44.1 
Yttrium,  Y            89.0 
Erbium,  Er         167.4 
Ytterbium,  Yt     172.0 

As  raw  material  for  the  rare-earth  preparations,  the  mixture  of  carbon- 
ates obtained  in  working  up  monazite  sand  (No.  170)  may  be  used.  The 
crude  cerium  carbonate,  or  cerium  oxalate,  that  can  be  obtained  on  the  market, 
furnishes  practically  the  same  mixture  of  the  earths,  since  it  usually  con- 
tains 40  to  50%  of  cerium  salt,  15  to  20%  of  lanthanum  salt,  25  to  30%  of 
didymium  salt,  and  5  to  8%  of  yttrium  earths. 

The  cerium  is  precipitated,  according  to  the  method  of  Witt  and  Theel,  by 
adding  ammonium  persulphate  to  a  boiling  solution  of  the  nitrates;  eerie  sul- 


SEPARATION  OF   THE  RARE  EARTHS.  239 

phate  is  thereby  formed,  but  this  is  partially  hydrolyzed,  and  an  insoluble  basic 
salt  precipitates.  By  neutralizing  the  acid  set  free  by  the  hydrolysis,  a  com- 
plete precipitation  is  made  possible,  but  on  the  other  hand,  it  is  necessary  to 
keep  the  liquid  slightly  acid  to  prevent  the  other  rare  earths  from  precipi- 
tating with  the  cerium.  All  of  the  cerium  is  thrown  down  in  this  way,  but 
it  is  contaminated  with  some  lanthanum  and  didymium. 

Lanthanum  and  didymium  are  thrown  out  of  the  nitrate  as  potassium 
double  sulphates;  the  yttrium  earths  remain  in  solution  and  can  be  precipi- 
tated as  oxalates  by  adding  ammonium  oxalate. 

The  products  of  this  separation  are  then  each  further  separated  and 
purified. 

Dissolve  100  g.  of  the  raw  material  by  warming  it  on  the  water 
bath  in  200  g.  of  concentrated  nitric  acid.  The  crude  carbonate 
dissolves  very  quickly,  but,  if  cerium  oxalate  is  used,  the  addi- 
tion of  fuming  nitric  acid  is  necessary  to  effect  solution.  Evapo- 
rate the  solution  on  the  water  bath  until  it  is  of  sirupy  consis- 
tency, and  then  take  it  up  in  1500  c.c.  of  water.  Add  35  g.  of 
ammonium  persulphate  to  this  solution  and  heat  it  to  boiling  in 
a  large  evaporating  dish.  Stir  the  liquid  by  means  of  a  mechan- 
ical stirrer,  and  add  powdered  magnesium  carbonate  in  small  por- 
tions to  the  boiling  mixture  until  finally  Congo  paper  is  no  longer 
turned  blue,  although  litmus  is  still  reddened  by  the  solution,  — 
toward  the  last  the  magnesium  carbonate  should  be  added  very 
cautiously.  The  reaction  is  complete  when,  even  after  boiling 
for  five  minutes,  the  liquid  does  not  become  acid  to  Congo  paper, 
although  it  must  still  turn  litmus  red.  ^  About  40  g.  of  the  mag- 
nesium carbonate  are  required.  Allow  the  dull-yellow  precipi- 
tate to  settle,  drain  it  on  the  suction  filter  while  it  is  still  warm, 
and  wash  it  with  hot  water. 

Test  the  filtrate  for  cerium  as  follows:  To  2  c.c.  of  the  solution 
add  ammonium  chloride  and  then  ammonia;  filter  off  and  wash 
the  precipitate  and  then  dissolve  it  in  5  to  10  c.c.  of  hot  concen- 
trated potassium  carbonate  solution.  Treat  the  solution  with 
hydrogen  peroxide  and  some  ammonia,  and  warm  it.  A  slimy 
precipitate  is  produced,  and  if  cerium  is  present  both  the  solu- 
tion and  the  supernatant  liquid  are  of  an  orange-yellow  color;  if 
cerium  is  absent  the  precipitate  is  white  or  a  faint  pink.  The 
test  is  less  sensitive  if  the  original  solution  is  warmed  directly  with 
hydrogen  peroxide  and  sodium  carbonate.  If  the  above  tests 
show  that  cerium  is  present,  add  more  ammonium  persulphate  to 


240  COMPOUNDS   OF  THE   RARE   ELEMENTS. 

the  solution  and  precipitate  the  cerium  as  before  by  adding  mag- 
nesium carbonate. 

Bring  the  cerium-free  nitrate  to  boiling  in  an  evaporating  dish, 
and,  while  stirring  as  before  with  the  mechanical  device,  add 
powdered  potassium  sulphate  (about  35  g.  in  all)  until  the  didy- 
miuni  absorption  bands  can  scarcely  be  detected  spectroscopically 
in  a  filtered  sample  of  the  solution.  In  making  this  test,  use  a 
pocket  spectroscope;  fill  a  test-tube  of  2  to  3  cm.  diameter  —  or 
still  better  a  parallel-walled  vessel  of  the  same  thickness  —  with 
the  solution  and  place  it  between  the  slit  of  the  spectroscope  and 
a  Welsbach  light.  The  most  easily  recognized  of  the  absorption 
bands  of  didymium  are  those  to  the  right  and  left  of  the  sodium 
line.  If  the  solution  is  almost  free  from  didymium,  drain  the  pre- 
cipitate of  the  double  sulphates,  M2(SO4)3  -3  K2SO4,  of  didymium 
and  lanthanum,  and  wash  it  with  a  dilute  potassium  sulphate 
solution. 

Add  ammonium  oxalate  to  the  filtrate,  and  after  some  hours 
wash  the  precipitate  with  cold  water;  this  precipitate  contains  the 
remainder  of  the  rare  earths  that  were  present  in  the  monazite  — • 
chiefly  the  yttrium  earths;  these  are  not  to  be  further  separated. 

(a)   Cerium  Compounds. 

Working  Up  of  the  Cerium  Precipitate. 

Cerium  may  be  readily  purified  from  other  metals  by  forming  eerie  ammo- 
nium nitrate,  (NH4)2Ce(NO3)6,  which  is  difficultly  soluble  in  nitric  acid. 

Weigh  the  dried  cerium  precipitate  approximately,  and  boil  it 
in  an  evaporating  dish  for  30  minutes  with  five  times  its  weight  of 
a  10%  solution  of  sodium  hydroxide.  After  letting  the  solution 
settle,  decant  off  the  liquid  as  completely  as  possible,  and  boil 
the  residue  with  some  fresh  caustic  soda  solution.  Pour  off  the 
solution  again,  and  wash  the  residue  by  decantation  with  hot 
water,  pouring  all  the  liquid  through  a  filter.  Finally,  collect  the 
entire  residue  on  the  filter,  wash  it  until  free  from  soluble  sul- 
phate, and  dry  it  in  the  hot  closet. 

Treat  the  eerie  hydroxide  thus  obtained  with  2.65  parts  by 
weight  of  concentrated  nitric  acid  (sp.  gr.  1.4),  filter  the  solu- 
tion through  an  asbestos  felt  in  a  Gooch  crucible,  and  add  to 
the  filtrate  a  hot  solution  of  0.39  parts  of  ammonium  nitrate  in 


CERIUM  COMPOUNDS  241 

1.15  parts  of  water.  Evaporate  the  solution  on  the  water  bath 
until  it  begins  to  crystallize,  then,  after  12  hours,  drain  the  red 
crystals  of  the  double  nitrate  and  wash  them  with  a  little  nitric 
acid.  Evaporate  the  mother-liquor  to  obtain  several  further 
fractions  of  crystals,  and  collect  each  fraction  by  itself.  From 
the  last  mother-liquor,  precipitate  the  remaining  cerium,  together 
with  the  impurities,  by  diluting  with  water  and  adding  oxalic 
acid. 

Ignite  a  small  portion  from  each  of  the  crystal  fractions  on  a 
porcelain  crucible  cover  over  the  blast  lamp.  The  residue  of 
eerie  oxide  is  pale  yellow  when  the  cerium  preparation  is  pure; 
if  didymium  is  present  it  is  of  a  reddish  to  chocolate-brown  color. 
Recrystallize  the  fractions  that  are  shown  to  be  impure  by  dissolv- 
ing the  double  nitrate  in  1.6  to  1.7  times  its  weight  of  40%  nitric 
acid  (3  parts  HNO3,  sp.  gr.  1.4,  and  2  parts  H2O). 

For  the  successful  preparation  of  eerie  ammonium  nitrate  it  is 
important  to  avoid  reducing  the  tetravalent  cerium.  On  this 
account  the  nitric  acid  which  is  used  as  solvent  must  be  boiled 
for  a  short  time  in  a  flask  in  order  to  expel  lower  oxides  of 
nitrogen. 

Cerium  Sulphides. 

Sulphides  of  cerium  can  be  prepared  by  igniting  cerous  sulphate  in  a 
stream  of  hydrogen  sulphide.  If  the  temperature  remains  below  720°  the 
dark-brown  disulphide  CeS2  is  formed;  but  if  it  is  kept  at  a  bright  red  heat, 
dark  cinnabar-red  cerous  sulphide,  C^Ss,  is  produced. 


Add  hydrogen  peroxide  to  a  boiling  aqueous  solution,  con- 
taining 20  g.  of  eerie  ammonium  nitrate,  until  the  liquid  is  decolor- 
ized; then  add  7  g.  of  sulphuric  acid  and  evaporate  the  solution 
to  dryness.  Place  the  residue  in  a  short  combustion  tube  and 
heat  it  in  a  stream  of  hydrogen  sulphide  which  has  been  dried 
with  calcium  chloride.  Heat  the  tube  in  a  combustion  furnace 
first  to  a  dull  red  and  later  to  a  bright  red,  and  occasionally 
revolve  the  tube  on  its  long  axis. 

If  cerous  sulphide  has  been  formed,  the  product  dissolves  in 
hydrochloric  acid  without  residue;  if  the  product  contains  cerium 
disulphide,  free  sulphur  separates  on  this  treatment.  Test  the 
hydrochloric  acid  solution  with  barium  chloride  to  show  whether 
the  sulphate  has  been  completely  decomposed. 


242  COMPOUNDS   OF   THE   RARE  ELEMENTS. 

Cerous  Chloride,  CeCl3. 

Anhydrous  cerous  chloride  is  readily  obtained  from  the  sulphide  by  ignit- 
ing the  latter  in  a  stream  of  hydrogen  chloride  gas. 

First  prepare  cerium  sulphide  in  the  manner  just  described; 
allow  the  tube  to  partly  cool,  and  pass  through  it  a  stream  of 
thoroughly  dry  hydrogen  chloride;  then  heat  it  again  to  dull  red- 
ness. The  transformation  is  complete  when  the  preparation  has 
become  pure  white. 

Anhydrous  cerous  chloride  forms  a  crystalline  very  hygroscopic 
mass. 

Anhydrous  Ceric  Sulphate,  Ce(S04}2.  Prepare  eerie  oxide  by 
igniting  eerie  ammonium  nitrate  strongly;  pulverize  the  oxide, 
and  boil  it  with  a  large  excess  of  concentrated  sulphuric  acid, 
whereby  it  is  converted  into  the  deep-yellow  crystalline  sulphate 
without  being  dissolved.  Pour  off  the  excess  of  acid,  wash  the 
residue  by  decantation  with  glacial  acetic  acid,  drain  it  on  a 
hardened  filter  paper,  and  dry  it  over  lime  in  a  vacuum  desiccator. 
Ceric  sulphate  dissolves  to  a  considerable  extent  but  never  com- 
pletely in  water;  on  boiling  the  dilute  solution  an  insoluble  basic 
salt  is  precipitated.  This  behavior  is  utilized,  as  has  been  shown, 
in  the  separation  of  cerium  from  the  other  rare  earths. 


(b)  Lanthanum  Compounds. 

Lanthanum  Carbonate  from  the  Lanthanum- Didymium  Precipi- 
tate. Boil  up  the  precipitate  of  the  potassium  double  sulphates 
with  five  parts  of  concentrated  nitric  acid  in  a  porcelain  casserole, 
and  pour  the  entire  mass  into  fifteen  parts  of  boiling  water.  To 
the  clear  solution,  which  is  colored  violet  by  neodymium,  add 
ammonium  oxalate,  — taking  0.7  to  0.8  g.  to  each  1.0  g.  of  the 
double  salt,  —  neutralize  the  solution  with  ammonia,  and  after 
some  time  filter  off  the  precipitate.  Wash  and  dry  the  latter,  and 
ignite  it  over  a  Fletcher  burner  in  a  porcelain  or  clay  crucible, 
surrounded  by  a  funnel  to  prevent  loss  of  heat.  Dissolve  the 
brown  oxide  thus  formed  by  heating  on  the  water  bath  with  as 
little  concentrated  nitric  acid  as  possible,  evaporate  the  solution 
to  a  sirupy  consistency,  and  dissolve  the  nitrates  in  1  liter  of 
water. 


LANTHANUM  COMPOUNDS.  243 

To  separate  the  didymium  from  the  lanthanum,  sift  magne- 
sium oxide  very  slowly  through  wire  gauze  into  the  solution  of 
the  nitrates,  while  this  is  kept  boiling  in  a  porcelain  dish  and  is  at 
the  same  time  stirred  with  a  mechanical  stirrer.  The  didymium 
is  precipitated  together  with  a  little  lanthanum,  but  the  greater 
part  of  the  lanthanum  remains  in  solution  free  from  didymium. 
Control  the  separation  by  means  of  the  spectroscope,  and  stop 
adding  magnesium  oxide  when  a  filtered  portion  of  the  solution, 
in  a  layer  3  cm.  thick,  shows  no  trace  of  the  didymium  absorp- 
tion bands. 

Filter  off  and  wash  the  precipitate  and  work  it  up  for  didy- 
mium as  directed  below  under  (c).  Precipitate  the  lanthanum 
from  the  filtrate  by  means  of  ammonium  carbonate  solution, 
after  first  adding  10  g.  of  ammonium  chloride.  In  order  to 
remove  all  traces  of  magnesium,  dissolve  the  precipitate  in  hydro- 
chloric acid  and  reprecipitate  it  with  ammonium  carbonate. 

Lanthanum  Sulphate,  La2(S04)3  •  9H20.  The  ennea-hydrate 
of  lanthanum  sulphate  is  readily  obtained  from  lanthanum  oxide, 
which  is  itself  prepared  by  igniting  the  carbonate.  Following 
the  directions  given  for  the  preparation  of  thorium  sulphate 
(No.  170)  treat  the  oxide  with  concentrated  sulphuric  acid,  dis- 
solve the  anhydrous  sulphate  in  ice-water,  and  then  warm  the 
solution. 

Determine  the  atomic  weight  of  lanthanum  by  the  sulphate 
method  (cf.  p.  236). 

Lanthanum  Acetate,  La(C2H302)3.3  H20.  Ignite  5  g.  of  lan- 
thanum carbonate  in  a  platinum  crucible,  pulverize  the  oxide  thus 
formed,  and  treat  it  in  a  flask  with  three  times  its  weight  of  gla- 
cial acetic  acid;  a  reaction  takes  place  with  strong  evolution  of 
heat,  either  of  itself  or  upon  gentle  heating.  Dissolve  the  result- 
ing thick  paste  in  as  little  water  as  possible,  whereby  only  a  very 
small  residue  remains,  and  evaporate  the  filtrate  until  crystalliza- 
tion takes  place.  Drain  the  crystals,  wash  them,  first  with  50% 
alcohol,  and  then  with  pure  alcohol,  and  dry  them  in  the  hot 
closet. 

Dependent  preparation:   Lanthanum  Blue,  No.  23. 

Lanthanum  Sulphide  and  Anhydrous  Lanthanum  Chloride  can 
be  prepared  according  to  the  methods  given  for  the  correspond- 
ing cerium  compounds. 


244  COMPOUNDS   OF  THE  RARE  ELEMENTS. 

(c)  Didymium  Compounds. 

Dissolve  the  didymium  precipitate  in  as  little  warm  nitric  acid 
as  possible,  and  treat  the  solution  with  magnesium  oxide,  as 
directed  in  the  didymium-lanthanum  separation,  until  the  didy- 
mium lines  are  only  just  faintly  visible.  Drain  and  wash  the 
resulting  precipitate,  dissolve  it  in  hydrochloric  acid,  and  repre- 
cipitate  the  didymium  with  ammonium  carbonate;  in  order  to 
remove  the  last  traces  of  magnesium,  dissolve  and  again  reprecipi- 
tate  the  didymium  carbonate.  The  preparation  thus  obtained 
consists  chiefly  of  the  neodymium  salt,  and  is  consequently  pink 
in  color.  A  separation  from  praseodymium  is  not  possible  on  a 
small  scale. 

Didymium  Chloride  with  Alcohol  ofCrystallization,DiCl3'3C2H6OH. 
Ignite  5  g.  of  didymium  carbonate  in  a  platinum  crucible,  dissolve 
the  powdered  didymium  oxide  by  boiling  with  a  saturated  solu- 
tion of  dry  hydrogen  chloride  in  anhydrous  alcohol  (about  50  c.c. 
are  necessary),  and  filter  the  yellow  solution,  which  has  the  con- 
sistency of  thin  sirup,  through  a  felt  of  asbestos  in  a  Gooch  cruci- 
ble. Saturate  the  filtrate  with  dry  hydrogen  chloride  gas.  Drain 
the  large  light-red  crystals  which  separate  in  the  course  of  several 
hours,  and  wash  them  with  a  little  alcohol.  Yield,  6  to  7  g. 

Didymium  Sulphide  and  Anhydrous  Didymium  Chloride  can  be 
prepared  according  to  the  methods  given  for  the  corresponding 
cerium  compounds. 


INDEX. 


PAGE 

Acetoneoxime 157,  158 

Acetone-semicarbazone 157,  161 

Acetylene 98 

Acids,  bases,  and  salts 58 

Addition  products  of  complex  hydro-metal-cyanic  acids 146 

Adsorbed  water 34, 188 

Adsorption 34,  38  et  seq. 

Adsorption  compounds ^  .  34 

curve 38 

equilibrium 38 

Affinity  of  condition 35 

Aggregation,  differences  in  the  state  of 15 

Alcosols 33 

Allotropy  of  cuprous  mercuriiodide 31 

of  mercuric  iodide 145 

of  mercuric  sulphide. , 88 

of  selenium 18 

of  silver  mercuriiodide 31 

of  silver  sulphide 28 

of  sulphur 29 

of  sulphur  trioxide 52 

Alum  from  kaolin 130 

Aluminium  amalgam 11 

boat  ..eo..      101 

chloride,  anhydrous 70 

heat  of  combustion  of 11 

sulphide.  .,.„.. .  „ 14 

Aluminothermy 11 

Aminomethanedisulphonic  acid. 158, 159 

Ammonia,  from  the  air  .'...». 95 

liquid . .    . .  i . .  .V* 182 

Ammonia  soda  process  ,..*« 134 

Ammonium  amalgam 20 

chloride,  dissociation  of 156 

copper  tetrasulphide 139 

cupric  sulphate 166 

pentasulphide 106 

phosphomolybdate ...... 151 

245 


246  INDEX. 

PAGE 

Ammonium  platinosulphite 184 

plumbic  chloride 104 

tribromide 107 

Ammonium  compounds 156 

Amorphous  state 32 

Amyl  nitrite 207 

Antimony 17, 19 

basic  chloride  of 17,  76 

pentachloride 76 

sulphate 129 

sulphide,  colloidal 37 

trichloride 75 

Apparatus  for  the  preparation  of  solid  chlorides 68,  78,  225,  226 

for  the  preparation  of  liquid  chlorides 73,  82 

Aquopentamminecobaltic  salts 172, 173, 176 

Arsenic  acid 55 

Asbestos  heating  box 3 

platinized 52 

Assaying 21 

Atomic  weight  determination  according  to  the  sulphate  method 237,  243 

Autooxidation 105 

Babo  boiling  funnel 3 

Barium  chloride  from  witherite 67 

Barium  chloride  dihydrate,  eutectic  point  of,  with  water 191 

Barium  dithionate 131 

ferrate 126 

hypophosphite 135 

nitrate  from  heavy  spar 129 

peroxide 102 

peroxide  hydrate 102 

trithiocarbonate 137 

Becquerel  rays 229 

Benzene,  from  acetylene 98 

Benzoic  acid,  distribution  of,  between  two  solvents 109 

Beryllium  acetate,  basic 220 

hydroxide '  219 

Bismuth  iodide,  basic 76 

nitrate 124 

nitrate,  basic 124 

tribromide 77 

triiodide 76 

Blast  furnace  equilibrium f 49 

Boiling-point  determination I 7 

Boiling-point  determination,  with  electrical  pyrometer 78 

Boron,  crystallized \. 14 

nitride. . ,  94 


INDEX.  247 

PAGE 

Bredig 36 

Bumping,  methods  of  avoiding 4 

Cadmium  iodide 148 

Calcium  hydride 57 

nitride 96 

sulphate,  hydrates  of 193 

Calibration  of  the  pyrometer 87 

of  the  spectroscope 218 

Camphor,  as  catalyzer 199 

Carbides 97 

Carbonatotetramminecobaltic  salts 171, 173 

Carbonyl  groups,  detection  of 157 

Cassius,  Purple  of 19,  35,  42 

Catalysis 36,  52,  62, 103, 199 

poisoning  of 37,  52 

Central  atom 141 

Cerium  compounds 240 

hydride 56 

Chamber  crystals 129,  204 

Changes  of  condition 26 

Chlorine,  preparation  of 69 

Chlorides,  preparation  of,  from  oxides 81 

Chlorination 199 

Chloroaquotetramminecobaltic  salts 172, 181 

Chloronitritotetramminecobaltic  salts 179 

Chloropentamminechromic  salts 182, 183 

Chloropentamminecobaltic  salts 172, 173, 174,  181 

Chloroplatinous  acid 184 

Chromic  chloride,  hydrates  of 197 

Chromic  oxide 55 

Chromium  metal 13 

nitride. 96 

trichloride,  anhydrous 71 

hydrates 197 

Cinnabar 20,  88 

cis-Position 181 

Clay  mantle 1 

Cobalt  acetate 145 

Cobalt  compounds  free  from  nickel 177 

Cobaltiammonia  compounds 167  et  seq. 

Cobaltous  salt  of  hydromercurithiocyanic  acid 143 

Colloidal  solution 33 

Colloidal  state 33 

Columbite 222 

Columbium  compounds 221,  222 

Complex  cations 155, 164, 165 


248  INDEX. 

PAGE 

Complex  compounds 99,  139,  155 

classification  of 100 

Complex  cyanogen  compounds 139 

Complexes,  formation  of 140 

stability  of 140 

Compounds,  simple 44 

containing  a  complex  negative  component 99 

containing  a  complex  positive  component 155 

Concentration,  anomalous  change  in 140 

Concentration,  molal 46, 110 

Condensation  of  very  volatile  substances 7,  8,  50 

Condensed  acids 151 

Conductivity,  molecular,  of  the  metal-ammonia  compounds 164, 165 

Contact  process 46,  51 

Convective  transference 35 

Cooling  curve 28,  30 

Coordination  number % 141,  165 

theory 141,  165 

Copper  ferrocyanide  membrane 42 

hydride 57 

sulphate 49 

Correction  of  melting  and  boiling  points 16,  79 

Coulomb 22 

Critical  point 51 

Croceocobaltic  salts 179 

Cryohydrates 187,  192 

Crystallization 5 

Crystalloids 33 

Crystals,  growth  of  large  at  expense  of  small  grains 174 

Cupellation 21 

Cupric  ammonium  sulphate 169 

bromide 65 

sulphate 49 

hydrates  of 188 

Cuprous  bromide 65 

chloride 66 

mercuriiodide 31 

oxide 56 

Curie,  M.  and  Mme 231 

Current  density 115 

yield 24, 116 

Cyanic  acid,  constitution  of 113 

Cyanogen 93,  94 

Cyanogen  compounds,  complex 139 

Decomposition  potential 59 

pressure ,• 163 


INDEX.  249 

PAGE 

Decomposition  temperature 163 

Density  determination  of  solid  substances 10,  24 

Deville,  H.  St.  Claire 103 

Devitrification  of  glass 33 

Dialysis 33,  37,  39 

Diazomethanedisulphonic  acid 158,  159 

Dibromopraseo  salts 1 78 

Dibromotetramminecobaltic  salts 178 

Dichlorodiammineplatinum 185 

Dichloropraseo  salts 180 

Dichlorotetramminecobaltic  salts 1 80 

Didymium  compounds 243 

Diethyl  sulphite,  symmetrical 207 

unsymmetrical 209 

Dilution  law 59 

Dimercuriammonium  hydroxide 162 

Dinitritotetramminecobaltic  salts 179,  180 

Diphenyl 215 

Diphenyl  lead  iodide 214 

Dissociation,  electrolytic 58,  100 

detection  of 61 

Dissociation  of  carbon  dioxide 10 

of  complex  compounds 100, 156 

Dissociation,  thermic 26,  45,  156 

Distillation 5 

fractional 203,  209,  211 

Dithio-oxamide 93 

Double  salts 99 

Efficiency  of  reactions 47 

Electroaffinity 59 

Electrode  surface 115 

Electrolysis 22,  23 

Electrolytes. 58 

precipitating  action  of,  on  colloids 34,  36 

Electrons 58 

Elements,  extraction  of 9 

occurrence  of 9 

rare 217 

Enantiotropy 27 

Endothermic  reactions 47,  92,  98, 103 

Equilibrium,  state  of 45 

in  blast  furnace 49 

Ester  equilibrium ., 206 

Esters 205 

of  nitric  acid 206 

of  nitrous  acid 207 


250  INDEX. 

PAGE 

Esters  of  phosphoric  acid 210 

of  silicic  acid 211 

of  sulphurous  acid 207,  209 

Ether,  compounds  of  complex  acids  with 146, 147, 151 

Ethyl  nitrate 206 

Eutectic  of  barium-chloride-dihydrate/water 192 

Eutectic  point 187, 191 

Evaporation 4 

of  inflammable  liquids 214 

Exothermic  reactions 47 

Faraday's  law 22 

Fehling's  solution 56 

Ferrate 117 

Ferric  chloride,  anhydrous 42 

Filter  paper,  hardened 4 

Filtration 4 

Fire-assay 21 

Fixed  points  in  thermometry 196 

Flavocobaltic  salts 178 

Fluorescein 54 

Fractional  distillation 203,  209,  211 

Fractionating  flasks 5,  6 

Frankland,  E 212 

Freezing  mixture 49, 183, 187 

Freezing-point  curve 187 

Furnaces 1,  2,  52 

Gold,  colloidal  solution  of 40 

extraction  of 18 

ruby  glass 40 

Goldschmidt  process 11 

Graham,  T 33 

Grignard's  reagent 213 

Guldberg,  C.  M 46 

Gypsum 193 

Halogen  compounds 65 

Halogens,  complex  acids  of 139 

complex  salts  of 139 

oxyacids  of 114 

Heating  curve 28 

of  crucibles 

of  dishes 4 

of  flasks • 

of  tubes 

with  a  reflux  condenser 3 


INDEX.  251 

PAGE 

Hexamminechromic  nitrate 182 

Hexamminecobaltic  salts 175 

Hexamminenickelous  bromide 171 

Hittorf,  W 140, 148 

Homogeneous  complexes 101 

Hot-cold  tubes 103 

Hydrates,  theory  of 186 

of  calcium  sulphate 193 

of  chromium  chloride 197 

of  sodium  sulphate 194 

Hydration  of  the  ions 189 

Hydrazine  sulphate 158 

Hydrides 56 

Hydrobromic  acid 62 

Hydrochloric  acid 70 

Hydrocobalticyanic  acid 147 

Hydrocyanic  acid ^ . .       92 

Hydroferrocyanic  acid 146, 147 

Hydrofluosilicic  acid 141 

Hydrogels 34 

Hydrogen  chloride,  preparation  of 70 

Hydrogen  cyanide 92 

Hydrogen  peroxide 103-106 

catalysis 36 

reactions 83, 104 

Hydrolysis 60 

Hydrosols 33 

Hydroxylaminedisulphonic  acid 156, 158 

Hydroxylamine  sulphate 156 

Hyposulphurous  acid 132 

Ice  curve 192 

Ignition  powder 12 

Indirect  combination 169 

Inner  sphere 165, 180 

Inoculation 30,  32, 194, 195 

lodic  acid 119 

anhydride 119 

Iodide-starch 39 

lonization  tendency 59 

Iron,  passive 32 

Irreversible  colloids 34 

Isomerism,  dynamic 208 

spacial 166, 180, 185 

Kaolin 130 

Kiihne's  dialyzing  tube 37 


252  INDEX. 

PAGE 

Lakes 40 

Lanthanum  blue 38 

compounds 242 

Lead  from  galena 20 

from  lead  oxide 10 

Lead  tetrachloride 143 

Le  Chatelier  pyrometer 86 

Lepidolite 217 

Lithium 22 

Lithium  carbonate  from  minerals 217 

Luteochromic  salt 182, 183 

Luteocobaltic  salts..  133 


sium  nitrate  hexahydrate 189 

nitride 95 

phosphide 97 

Magnus,  salt  of 184 

Manganates 124 

Manganese 12 

Manganese  chloride 68 

sulphide 89 

Mangano-manganic  oxide 12 

Mantle,  clay 1 

Marignac , 221 

Mass  action  law 46,  59,  60, 109,  206 

Maxima,  concealed 189 

Melting-point,  correction  of 16,  79 

determination 15 

diagram 191 

Mercuric  cyanide 92 

iodide .-     31,  144 

oxide 162 

sulphide 88 

Mercury  from  cinnabar 20 

Metal-ammonia  compounds 163 

Metals,  base  and  noble 60 

reduction  by  electrolysis 22 

Metal-organic  compounds 211 

Methylnitrite 207 

Millon's  base 162 

salts  of 162 

Molybdenum  blue 39 

compounds 222 

Molybdenum,  metal 224 

Monazite 235 

Monochlorsulphuric  acid 199,  201 

Monochloramine 158, 160 


INDEX.  253 

PAGE 

Monotropy 27 

Mosaic  gold 89 

Neutralization i 60 

Neutron 59 

Nickel,  finely  divided 215 

cobalt-free 171 

Nickel  air  bath 3 

carbonyl. 215 

oxalate 215 

Nitrate  method 123 

Nitrates  and  nitrites 121 

Nitric  oxide-metal  compounds 155 

Nitrides 92 

Nitrilosulphonate  of  sodium 156 

Nitrito-acids 149 

Nitrogen  dioxide ._. .  54,  55 

trioxide 55 

Nitrososulphonates 156 

Nitrosylsulphuric  acid 129,  204 

Nitrous  oxide 112 

Non-electrolytes,  complex 199 

Oersted 81 

Optical  rotation 153 

Organo-complexes 152 

Osmotic  pressure 42,  43 

Ostwald's  dilution  law f 59 

Oxide  hydrogels 34 

Oxides 49 

Oxides,  hydrated 34 

Oxydimercurammonium  hydroxide 162,  163 

Oxycobaltammine  chloride 173 

Palladium-hydrogen 56 

Passive  condition 32 

Peptonization 34 

Perchloric  acid 118 

Peroxides 101 

Petalite 217 

Peyrone,  Salt  of 185 

Phosphides 97 

Phosphorus,  constitution  of  acids  of 135 

oxychloride  75 

pentachloride 74 

as  a  chlorinating  agent 199,  201 

pentasulphide 85 


254  INDEX. 

PAGE 

Phosphorus  residues,  disposal  of 74 

trichloride 74 

Phosphine 97 

Phthalic  acid  from  naphthaline 53 

acid  anhydride 54 

Pitchblende 229 

Platinized  asbestos 52 

Platino-ammonia  compounds 168, 184 

Platinotypes 153 

Platinum  ammonia  compounds 166  et  seq. 

as  catalyzer 36,  52,  62 

colloidal 36 

Platosammine  chloride 185 

Poisoning  of  catalyzers 37 

Polyhalogen  compounds 107 

Polymerization 26 

Polynitrides Ill 

Polysulphides 106 

Potassium  bromate 121 

bromide 121 

chlorate 117 

cobalticyanide 145 

cobaltinitrite 150 

cobaltothiocyanate .' 147 

columbium  oxyfluoride 221 

cyanate 113 

didymium  sulphate 240 

ferric  oxalate : 152 

ferric  sulphide 138 

iodate  from  potassium  chlorate 120 

iodide 66 

lanthanum  sulphate 240,  243 

lead  iodide 144 

mercuriiodide 144 

mercurinitrite 149 

nitrate  from  sodium  nitrate 122 

perchlorate , 118 

permanganate 124-126 

persulphate,  electrolytically 133 

tantalum  heptafluoride 221 

tribromide 108, 109- 

triiodide 110 

trithiocarbonate  solution 136 

Potential  series 59 

Praseo  salts 178, 181 

Precipitating  colloids 40 

Precipitation  processes 19 


. 

INDEX.  255 

PAGE 

Protective  colloids 35,  40 

Prussic  acid  residues 93 

Pseudo  solutions 34 

Pulverizing 8 

Pure  compounds 45,  46 

Purple  of  Cassius 19,  35,  42 

Pyrometer 86 

Pyrosulphuric  acid  dichloride 202 

Radioactive  constituents  of  pitchblende 229-232 

Radioactivity 229,  231 

detection  of 231 

Rare  earths 237 

historical 237 

separation 239  et  seq. 

Reactivity  and  degree  of  dissociation 45 

Reduction  with  aluminium "...        11 

with  aqueous  reducing  agents 17, 19,  40 

with  carbon 10 

with  carbon  monoxide 10,  41 

with  hydrogen. 84 

with  potassium  cyanide 15 

Reversible  colloids 34 

reactions 45, 102 

Richards,  T.  W 171, 196 

Rieset,  Base  of 186 

Roasting  processes 20 

Roseocobaltic  salts 176 

Rotation,  optical 153 

Row  burner 2 

Rubidium,  spectroscopic  detection  of 218 

iodide  tetrachloride Ill 

triiodide Ill 

Rutile 81 

Salting  out 34,  39 

Saponification 199 

of  esters 205  et  seq. 

Saturation  capacity  of  the  metals 212 

Schlippe's  salt •'. 137 

Sealing  flasks 8 

tubes 8,  50 

Secondary  valences 141, 155 

Selenium. 17 

dioxide 17 

Semicarbazide  hydrochloride 161 

Semipermeable  membranes 42 


256  INDEX. 

PAGE 

Silicates,  of  the  heavy  metals,  membranes  of 43 

Silicic  acid  hydrogel 151 

Silicon  chloroform 81 

crystallized 13 

hexachloride 81 

Silicon  octochloride 81 

tetrachloride 80 

Silicotungstic  acid 151 

Silver  from  silver  iodide 209 

by  cupellation 21 

Silver-ammonia  ion 164 

sulphate 169 

Silver  hydrazoate 112 

mercuriiodide 31 

sulphate 169 

sulphide 28 

sulphite 209 

Simple  compounds 44 

Sodamide .      Ill,  112 

Sodium  amalgam 20 

bicarbonate 134 

carbonate 134 

cobaltinitrite 149 

hydrazoate Ill,  112 

hypochlorite 114 

nitrite 121 

peroxide 101 

tetrathionate 132 

thioantimonate 137 

thiosulphate 131 

Solubility  curve 192 

Solvay  process 134 

Spectrum  analysis 217,  218 

Spodumene ' 217 

Stannic  chloride 79 

oxide,  colloidal 42 

sulphide 89 

Stannous  sulphide 88 

Sublimation 17 

Substituted  ammonium  compounds 156 

Succinic  acid,  distribution  of,  between  two  solvents 108 

Suction 5 

Sulphate  method  for  atomic  weight  determinations 237 

Sulphides 85 

Sulphur,  amorphous 33 

chloride,  S2C12 72 

constitution  of  oxy-acids  of 126 


INDEX.  257 

PAGE 

Sulphur  dioxide 49 

preparation  of 49,  71 

livers  of 106 

monoclinic 29 

rhombic 29 

Sulphuric  acid  from  pyrite 128 

by  the  contact  process 46,  51 

dichloride 199 

monochloride 199,  201 

Sulphurous  acid  chloride 202 

Sulphur  trioxide 51 

Sulphuryl  chloride 199 

Supercooling 16,  32 

Supersaturated  solutions 194,  195 

Tammann,  G  . 32,  188, 189 

Tantalum  compounds 221 

Temperature,  influence  of,  on  equilibrium  constants 47 

Tetrammine  cupric  sulphate 169 

platinous  chloride 185 

Tetranitrito-diammine-cobaltate  of  potassium 150 

Tetraphenyl  lead 213 

Thallic  oxide,  hydrated 64 

Thallous  hydroxide 64 

nitrate 64 

sulphate 64 

Thermic  analysis 187, 189 

Thermite 11 

Thermoelectric  temperature  measurements 86 

Thermometry,  fixed  points  in 196 

Thionyl  chloride 202 

Thorium,  atomic  weight  determination  of 237 

compounds 235 

Time  reaction 119 

Tin  from  cassiterite 15 

sulphides  of 88,  89 

tetrachloride 79 

Titanium  dioxide 83 

disulphide 90 

tetrachloride  from  rutile 81 

trichloride 84 

Transition  point,  determination  of 27,  31 

of  sodium  sulphate 195 

Zraris-Position 180 

Trinitritotriamminecobalt 182 

Tungsten  compounds 227 

Tyndall  phenomenon 41 


258  INDEX. 

PAGE 

Ultramicroscopic  investigation „ 34 

Uranium  compounds 232 

Uranyl  laevo-malate 153 

Urea 114 

Valencies,  principal 141 

van't  Hoff,  J.  H 48 

Vapor  tension  analysis 163, 188 

Waage,  P 40 

Washing  precipitates 5, 136 

Water,  detection  of  traces  of 144 

of  constitution 189 

Werner,  A 141, 156-169, 180 

Williamson's  violet 93 

Winkler,  Cl 51 

Wohler,  F 55,  96 

Yttrium  earths 237,  240 

Ziervogel  process 123 

Zinc  ethyl 212 

Zircon  gold  purple 42 

oxide,  colloidal 40 


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